so 


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LIBRARY 

OF   THE 

UNIVERSITY  OF  CALIFORNIA 

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Received 
Accession  No  403$  3  1-       Class  No. 


LABORATORY    WORK 


IN 


CHEMISTRY 


A  SERIES  OF  EXPERIMENTS  'IN  GENERAL 
INORGANIC  CHEMISTRY 


BY 

EDWARD    H.    REISER 

\^ 

PROFESSOR  OF  CHEMISTRY,  BRYN  MAWR  COLLEGE 


UNI7BESITY 


NEW  YORK-:- CINCINNATI-:. CHICAGO 

AMERICAN    BOOK    COMPANY 


COPYRIGHT,  1895,  BY 
AMERICAN  BOOK  COMPANY. 


LAB.  WORK  IN  CHEM. 
M 


or 

U5IVBRSITT 


PREFACE. 


THE  directions  for  laboratory  experiments  contained  in 
this  book  have  been  arranged  for  the  use  of  students  that 
are  following  a  course  of  lectures  or  recitations  on  general 
chemistry.  These  laboratory  exercises  are  intended  to 
illustrate  and  to  be  supplementary  to  the  work  of  the  class 
room.  It  is  not  intended  that  this  book  shall  take  the 
place  of  the  personal  instruction  of  the  teacher ;  on  the 
contrary,  its  object  is  to  facilitate  this  work  of  the  instruc- 
tor where  the  number  of  students  in  the  class  is  large. 

The  laboratory  work  and  the  lectures  or  recitations 
should  go  hand  in  hand.  Before  a  subject  is  taken  up  in 
the  laboratory  it  ought  to  be  briefly  outlined  in  the  class 
room,  and  the  object  of  the  experiments  that  are  to  be 
made  ought  to  have  been  clearly  brought  into  the  minds 
of  the  students.  At  the  same  time  they  must  not  be  told 
everything  that  they  are  expected  to  observe ;  they  must 
be  given  opportunity  to  acquire  the  art  of  observing  accu- 
rately and  of  describing  what  they  see.  After  all  the 
experiments  on  a  given  subject  have  been  made,  a  full  dis- 
cussion of  the  results  is  in  place.  Then  the  instructor 
should  see  to  it  that  the  students  have  observed  everything 
that  was  to  be  seen,  that  proper  conclusions  have  been 
drawn  from  the  results  of  the  experiments,  and  that  the 
connections  between  these  results  and  the  general  laws 
and  principles  of  the  science  are  clearly  brought  out. 

iii 


IV  PREFACE. 

Some  of  the  experiments  illustrating  important  facts 
and  laws  of  chemistry  are,  on  account  of  the  expensive 
apparatus  or  chemicals  required,  or  because  of  the  skill  and 
care  necessary  in  their  execution,  of  such  a  nature  that 
they  cannot  well  be  done  by  each  individual  member  of  a 
large  class ;  nevertheless  they  are  of  such  importance  that 
each  student  ought  to  have  an  accurate  knowledge  of  them. 
Experiments  of  this  kind  have  been  called  Laboratory 
Demonstrations  in  the  text ;  and  it  is  intended  that  these 
experiments  shall  be  carried  out  in  the  presence  of  the 
whole  class  by  one  or  two  of  the  more  skillful  students, 
working  under  the  immediate  supervision  of  the  instructor. 
After  each  such  demonstration,  opportunity  ought  to  be 
given  to  the  class  to  ask  questions  and  to  discuss  the 
results. 

In  doing  the  work  outlined  in  these  pages,  the  student  is 
expected  to  keep  a  full  record  of  each  experiment  (labora- 
tory demonstrations  included)  in  a  notebook  and  to  con- 
scientiously do  the  work  indicated.  The  notebook  ought 
to  be  examined  and  corrected  from  time  to  time  by  the 
instructor.  Each  student  should  own  a  good  text-book  of 
chemistry  and  ought  to  have  ready  access  to  the  chemical 
dictionaries  and  works  of  reference. 

In  conclusion  I  desire  to  express  my  thanks  to  Dr.  E.  P. 
Kohler  for  valuable  suggestions  made  to  me  in  regard  to 
the  directions  for  some  of  the  experiments,  and  also  to 
Professor  Ira  Remsen  for  permission  to  use  the  directions 
for  several  experiments  that  are  given  in  his  text-books. 

I  am  also  under  obligations  to  the  American  Book  Co., 
publishers  of  Storer  and  Lindsay's  Manual  of  Chemistry, 
for  the  privilege  of  using  some  of  the  illustrations  con- 
tained in  that  text-book. 


SUGGESTIONS   TO   STUDENTS 


BEFORE  beginning  an  experiment,  read  carefully  all  the 
directions  for  the  work.  Arrange  the  apparatus  as  directed, 
or  if  you  desire  to  use  a  modified  form  of  apparatus,  consult 
the  instructor  before  proceeding.  In  your  notebook  de- 
scribe the  apparatus  and  materials  that  you  use  and  what 
you  have  done  with  them.  A  simple  sketch  of  apparatus 
is  very  useful.  Describe  all  the  changes  that  occur,  and 
answer  the  questions  that  are  given  in  the  text. 

In  generating  gases  make  sure  that  all  joints  and  stop- 
pers of  the  apparatus  are  tight.  If  the  generating  vessels 
are  to  be  heated,  it  is  necessary  to  use  thin-walled  flasks. 
Thick  glass  bottles  almost  always  crack  when  heated. 
Small  quantities  of  liquids  are  heated  in  test  tubes.  Solids 
are  heated  in  ignition  tubes,  and  all  glassware  must  be  dry 
on  the  outside  before  heating.  Always  keep  your  appara- 
tus and  desk  clean  and  neat. 


TABLE   OF   CONTENTS. 


PAGE 

Exercises  in  the  Construction  of  Apparatus i 

Changes  in  Matter,  Chemical  and  Physical  Changes 3 

Matter  and  Energy 7 

Mechanical  Mixtures  and  Chemical  Compounds 8 

Homogeneous  Matter,  Elements,  Compounds,  Physical  Mixtures    .  9 

Chemical  Balance 10 

Measurement  of  Liquids io 

The  Air 11 

Measurement  of  Gases 14 

Oxygen 16 

The  Conservation  of  Matter 20 

Laws  of  Conservation  of  Matter  and  of  Definite  Proportions      .     .  21 

Nitrogen 22 

Volumetric  Composition  of  Air 23 

Water 23 

Hydrogen 26 

Composition  of  Water  by  Volume  and  by  Weight 32 

Ozone 36 

Hydrogen  Dioxide  and  the  Law  of  Multiple  Proportions       .     .     .  37 

Ammonia 39 

Nitric  Acid 40 

Nitrous  Oxide 43 

Nitric  Oxide 44 

Volumetric  Composition  of  Nitrous  and  Nitric  Oxides  and  the 

Law  of  Volumes 46 

Chlorine 47 

Hydrochloric  Acid 49 

Acids,  Bases,  and  Salts 52 

Carbon 54 

Marsh  Gas ' 55 

Ethylene 56 

Acetylene 56 

Carbon  Dioxide 57 

Carbon  Monoxide 59 

Influence  of  Temperature  upon  the  Rapidity  of  Chemical  Action    .  60 

Illuminating  Gas 61 

The  Blowpipe,  Oxidation,  and  Reduction 61 

The  Equivalent  Weights  of  the  Elements,  and  the  Law  of  Recip- 
rocal Proportions 62 

vii 


viii  CONTENTS. 

PAGE 

Molecular  and  Atomic  Weights      ...........  63 

Bromine 64 

Hydrobromic  Acid • 64 

Preparation  of  Ammonium  Bromide 65 

Preparation  of  Potassium  Bromide 66 

Iodine 66 

Hydriodic  Acid 67 

Fluorine 68 

Sulphur 69 

Hydrogen  Sulphide 70 

Sulphur  Dioxide 71 

Sulphuric  Acid 73 

Phosphorus 75 

Arsenic 77 

Antimony 79 

Boron 80 

Silicon 81 

Classification  of  the  Elements  and  the  Periodic  Law 83 

Potassium 83 

Preparation  of  Potassium  Iodide    . 86 

Preparation  of  Potassium  Chlorate 87 

Sodium 87 

The  Spectroscope 89 

Ammonium  Salts 89 

Volumetric  Composition  of  Ammonia 91 

Calcium     . 91 

Determination  of  Carbon  Dioxide  in  Calcium  Carbonate     ...  93 

Barium  and  Strontium 94 

Magnesium 95 

Zinc 96 

Cadmium 97 

Mercury 97 

Copper 99 

Silver 100 

Aluminium 102 

Tin 104 

Lead 105 

The  Law  of  Specific  Heats 106 

Bismuth 106 

Chromium 107 

Manganese 109 

Iron in 

Nickel  and  Cobalt  ..11 113 

Table  of  the  Elements 114 

Table  of  the  Weights  of  Gases 115 

Table  of  Tension  of  Aqueous  Vapor 116 

Index 117 


WORK    IN   CHEMISTRY. 


EXERCISES  IN  THE  CONSTRUCTION  OF  APPARATUS,  ETC. 

1.  A  Wash  Bottle.  —  Examine  carefully  the  construction 
of  the  sample  wash  bottle.  Soften  a  cork  by  means  of  the 
cork  press,  or  by  rolling  it  under  the  foot,  and  fit  it  air- 
tight into  the  neck  of  a  flask  of  about  one  liter  capacity. 
Cut  the  glass  tubing  to  the  proper  length  with  the  triangu- 
lar file,  and  bend  the  pieces  by  means  of  the  broad-top 
burner.  Round  the  ends  of  each  tube  by  holding  it  for  a 
short  time  in  the  flame.  Both  in  bending  and  rounding 
the  glass,  always  keep  turning  it  while  in  the  flame  so  that 
it  will  heat  uniformly.  Select  a  cork  borer  slightly  smaller 
in  diameter  than  the  glass  tubing,  and  carefully  bore  two 
parallel  holes  through  the  cork.  Enlarge  and  smooth  the 
holes  with  the  round  file,  and  fit  in  tightly  the  glass  tubes. 
Contract  the  outside  opening  of  the  glass  tube  which 
extends  to  the  bottom  of  the  bottle.  This  can  be  done 
either  by  drawing  out  the  end  to  a  small  diameter  and 
cutting  it  off  with  the  file,  or  by  rounding  the  end  in  the 
flame  until  the  opening  is  not  larger  than  the  diameter  of 
a  pin.  Clean  out  the  flask  and  tubes  thoroughly,  rinse 
with  distilled  water,  and  fill  the  flask  with  distilled  water. 


2  LABORATORY  WORK  IN  CHEMISTRY. 

2.  Glass  Stirring   Rods.  —  Cut  a  glass  rod  into  short 
pieces  that  are  of  the  proper  length  for  stirring  liquids  in 
the  beakers  with  which  you  are  provided,  and  round  both 
ends  of  each  piece  in  the  flame  of  the  burner. 

3.  Preparation  of  Cut  Filters.  —  Filter  paper  in  sheets, 
a  piece  of  cardboard,  and  a  pair  of  compasses  are  required. 
A  filter  when  folded  and  fitted  into  the  funnel  must  not 
quite  reach  the  edge  of  the  funnel.     Make  circles  on  the 
cardboard,  the  radii  of  which  are  equal  to  the  length  of  the 
sides  of  the  different  funnels.     Cut  out  these  circles  and 
use  them  as  patterns  in  cutting  the  sheets  of  filter  paper. 

4.  Folding   Filters.  —  Practice   folding   the   cut   filters 
(see   Fig.    i)   and   fitting   them    into   the   funnels.     After 

having  succeeded  with  the  plain  filter,  inquire 
of  the  instructor  how  star  filters  are  made. 
Practice  pouring  water  into  the  filter  from  a 
beaker  by  means  of  a  glass  rod. 


5.  Exercise  in  Manipulation.  Crystalliza- 
tion. —  Pulverize  a  spoonful  of  crystals  of 
FIG-  *•  lead  nitrate  with  the  mortar  and  pestle,  and 
pour  the  powder  on  a  piece  of  paper.  In  a  beaker,  heat 
water  to  boiling,  and  pour  into  it,  from  time  to  time,  some 
of  the  powder.  After  each  addition  of  the  powder,  stir 
with  a  glass  rod  until  all  is  dissolved.  Filter  the  hot  solu- 
tion and  evaporate  the  clear  filtrate  in  an  evaporating  dish. 
Set  aside  the  concentrated  solution  to  crystallize.  Repeat 
the  experiment,  using  copper  sulphate  and  potassium 
bichromate  instead  of  lead  nitrate. 

6.    Filtration  and  Washing  of  Precipitates.  —  To  a  clear 
solution  of  lead  nitrate  add  one  of  potassium  bichromate ; 


CHANGES  IN  MATTER.  3 

mix  thoroughly  with  a  glass  rod.  A  yellow  insoluble  pre- 
cipitate, called  chrome  yellow,  is  formed.  Let  the  mixture 
stand  for  a  few  minutes  until  the  precipitate  has  settled ; 
then,  by  means  of  the  glass  rod,  carefully  pour  the  clear 
liquid  into  a  filter.  When  nearly  all  has  been  decanted,  add 
distilled  water  to  the  precipitate  and  stir  with  the  rod. 
Now  allow  the  precipitate  to  subside,  and  again  decant  the 
clear  liquid.  Repeat  this  washing  with  water  two  or  three 
times,  or  until  a  drop  of  the  filtrate  evaporated  on  a  watch 
glass  leaves  no  residue.  Transfer  the  precipitate  to  the 
filter  by  means  of  a  jet  of  water  from  the  wash  bottle, 
allow  it  to  drain,  and  set  it  aside  to  dry. 


CHANGES  IN  MATTER. 

7.  By  means  of  pincers  hold  a  piece  of  platinum  wire  or 
foil  in  the  blue  flame  of  the  Bunsen  burner.     Notice  how 
the  platinum  is  changed.     Take  it  out  of  the  flame.     Has 
the  platinum  been  changed  permanently  ?     [Write  descrip- 
tions of  the  experiments  and  answers  to  the  questions  in 
your  laboratory  notebook.] 

8.  Heat  a  short  piece  (two  or  three  inches  long)  of  mag- 
nesium ribbon  in  the  same  way  in  which  you  heated  the 
platinum  wire.      What  happens  in  this    case  ?      Has  the 
magnesium  been  changed  permanently?     Has  a  new  sub- 
stance been  formed  ? 

9.  Into  a  clean,  dry  test  tube  put  enough  white  sugar  to 
make  a  layer  -^  of  an  inch  thick.     Hold  the  tube  in  the 
Bunsen  flame.     What  changes  take  place  ?     WThat  do  you 
notice  on  the  sides  of  the  tube  ?     What  remains  behind  ? 
What  is  its  color  and  taste  ?     Does  it  dissolve  in  water  ? 


4  LABORATORY   WORK  IN  CHEMISTRY. 

Is  it  sugar?     Has  the  sugar  been  changed  permanently 
into  other  substances  ? 

10.  Dissolve  a  small  quantity  (two  or  three  grams)*  of 
common  salt  in  water.     Evaporate  the  solution  to  dryness 
in  an  evaporating  dish.    Examine  the  residue.     Has  the  salt 
been  changed  permanently  when  it  has  gone  into  solution  ? 

11.  Into  a  small  tube  of  hard  glass  (ignition  tube)  put 
enough  red  oxide  of  mercury  to  form  a  layer  J  of  an  inch 
in  thickness.     Heat  the  tube.     What  change  in  color  do 
you  notice  ?     What  is  deposited  on  the  sides  of  the  tube  ? 
Insert  a  piece  of  wood  with  a  spark  on  the  end  into  the 
tube ;  keep  on  heating.     What  follows  ?     Take  it  out  and 
put  it  back  a  few  times.     Is  there  any  difference  between 
the  burning  in  the  tube  and  out  of  it  ?     What  difference  ? 
Keep  on  heating  until  nothing  is  left  on  the  bottom  of  the 
tube.     How  do  you  know  that  the  red  substance  that  you 
put  into  the  tube  has  been  changed  ?     Has  there  been  a 
material  alteration  in  composition  ? 

12.  Boil  water  in  a  beaker.     Bring  a  cold  surface  into 
the  steam.     What  collects  on  the  cold  surface  ?     Is  there 
any  change  of  composition  when  water  is  converted  into 
steam  ? 

In  which  of  the  preceding  experiments  have  the  sub- 
stances been  changed  permanently  so  that  new  substances 
having  different  properties  were  formed  ?  In  which  have 
the  changes  been  temporary  ?  What  are  changes  of  the 
first  kind  called  ?  What  are  those  of  the  second  kind 
called  ? 

Read  in  your  text-book  about  the  difference  between 
chemical  and  physical  changes. 


CHANGES   IN   MATTER.  5 

In  three  of  the  preceding  experiments  it  was  neces- 
sary to  heat  the  substances  in  order  that  the  chemical 
changes  might  take  place.  Under  certain  conditions 
light  and  electricity  bring  about  chemical  action.  Chemi- 
cal action,  too,  can  result  in  the  production  of  heat,  light, 
and  electricity.  Frequently  chemical  changes  take  place 
by  merely  bringing  substances  into  contact  with  one 
another. 

13.  Examine  a  piece  of  calc-spar  or  marble.  Notice 
whether  it  is  hard  or  soft.  Heat  a  small  piece  in  an  igni- 
tion tube.  Does  it  melt  or  change  in  any  way  ?  Is  the 
calc-spar  soluble  in  water  ?  In  order  to  learn  whether  a 
substance  is  soluble  in  water  proceed  as  follows  :  Put  a 
small  quantity  of  the  powdered  substance  in  a  test  tube 
with  distilled  water.  Shake  thoroughly,  and  then,  as  heat- 
ing usually  aids  solution,  boil.  Filter,  and  evaporate  a  few 
drops  of  the  clear  nitrate  in  an  evaporating  dish  or  watch 
glass.  If  there  is  anything  solid  in  solution,  there  will  be 
something  solid  left  in  the  dish  or  the  watch  glass.  If 
not,  there  will  be  nothing  left.  Knowing  now  the  general 
properties  of  the  calc-spar,  you  will  be  able  to  determine 
whether  it  is  changed  or  not.  Treat  a  small  piece  in  a 
test  tube  with  dilute  hydrochloric  acid.  What  takes  place  ? 
After  the  action  has  continued  for  about  a  minute,  insert  a 
lighted  match  in  the  upper  part  of  the  tube.  Does  the 
match  continue  to  burn  ?  Does  the  gas  in  the  tube  burn  ? 
Is  the  invisible  substance  in  the  upper  part  of  tube  ordi- 
nary air  ?  Why  not  ?  Does  the  calc-spar  disappear  ?  In 
order  to  tell  whether  it  has  changed  chemically,  the  solu- 
tion must  be  evaporated.  Pour  it  into  a  clean  evaporating 
dish  and  heat  with  a  small  flame.  Heat  gently  at  the  end, 
and  stir  with  a  glass  rod  to  avoid  spattering.  Examine 


6  LABORATORY   WORK   IN  CHEMISTRY. 

the  dry  substance  left  in  the  dish  and  compare  its  proper- 
ties with  those  of  the  substance  which  was  put  into  the 
test  tube.  Is  it  the  same  substance?  Is  it  hard  or  soft? 
Does  it  change  when  heated  in  a  tube  ?  Is  there  an 
appearance  of  bubbling  when  hydrochloric  acid  is  poured 
upon  it  ?  Does  it  change  when  allowed  to  lie  in  contact 
with  the  air  ?  When  calc-spar  or  marble  is  crushed  and 
ground  to  a  fine  powder,  is  it  changed  chemically  ? 
What  does  this  experiment  illustrate  ?' 

14.  Bring  together  in  a  test  tube  a  small  strip  of  copper 
and  some  strong  nitric  acid.     Hold  the  mouth  of  the  tube 
away  from  your  face  and  do  not  inhale  the  vapors.     What 
is  the  appearance  of  the  gas  given  off  ?    What  is  the  color 
of  the  liquid  in   the   tube  ?     Does   the  copper   dissolve  ? 
Pour  the  solution  into  an  evaporating  dish  and  evaporate 
to  dryness.     (Conduct  this  evaporation  in  the  hood,  as  the 
vapors  of  nitric  acid  are  annoying  in  the  air  of  the  room.) 
In  evaporating  to  dryness  be  careful  not  to  heat  the  residue 
to  too  high  a  temperature.     Describe  the  appearance  of 
the  substance  into  which  the  copper  has  been  changed. 
Is  it  soluble  in  water  ?     What  is  the  effect  of  heat  upon 
it  ?     When  copper  is  heated  so  that  it  melts,  is  the  change 
chemical  ? 

15.  Try  the  action  of  dilute  sulphuric  acid  on  a  little 
zinc  in  a  test  tube.     An  invisible  gas  will  be  given  off. 
Apply  a  light  to  the    mouth   of   the  tube.     What  takes 
place  ?     Add  more  zinc  to  the  acid,  and  when  it  no  longer 
acts  on  the  metal,  filter  some  of  the  solution  and  evaporate 
it   to  dryness.     Carefully  compare  the  properties  of  the 
substance  left  behind  with  those  of  the  zinc.     When  an 
electric  current  is  passed  through  a  strip  of  zinc,  and  it 
acquires  new  properties,  is  it  a  chemical  or  physical  change  ? 


MATTER  AND   ENERGY.  7 

16.  Pour  concentrated  nitric  acid  upon  a  bit  of  tin  in 
a   test   tube.       If    no  change   takes   place   at  first,  heat 
gently,  and  presently  you  will  have  evidence  that  change 
is  taking  place.     Is   there   anything   in  this  experiment 
which  suggests  Experiment  14  ?    What  is  left  behind  after 
the  action  is  finished  ? 

What  does  this  experiment  show  ? 

As  a  rule,  chemical  changes  take  place  more  readily  in 
gas  mixtures  and  in  solutions  than  between  solids. 

17.  Mix  together  in  a  dry  mortar  a  little  dry  tartaric 
acid  and  about  an  equal  quantity  of  dry  sodium  bicarbonate. 
Do  you  see  any  evidence  of  action  ?     Now  dissolve  a  little 
tartaric  acid  in  water  in  a  test  tube,  and  a  little  sodium 
bicarbonate  in  water  in  another  test  tube.     Pour  the  two 
solutions  together.      What  evidence   have  you    now  that 
action  is  taking  place  ?     Pour  water  upon  the  dry  mixture 
first  made.     Does  action  take  place  ?     What  causes   the 
bubbling  ?     Will   a   match  burn  in   the  gas  ?      In  which 
experiment  already  performed  was  a  similar  gas  obtained  ? 

In  which  of  the  previous  experiments  have  substances 
been  dissolved  in  liquids  without  undergoing  chemical 
change  ?  In  which  has  solution  been  accompanied  by 
chemical  action  ? 

MATTER  AND  ENERGY. 

Read  in  the  text-books  on  physics  and  chemistry  that 
are  accessible  to  you  about  the  relations  of  matter  and 
energy.  Can  matter  of  itself  undergo  change  ?  What  is 
it  that  causes  changes  in  matter  ?  What  are  the  different 
kinds  or  forms  of  energy  ?  What  is  the  characteristic 
property  of  energy  ?  Do  changes  in  energy  accompany 
changes  in  composition  of  bodies  ?  Give  illustrations. 


8  LABORATORY   WORK  IN  CHEMISTRY. 

MECHANICAL  MIXTURES   AND  CHEMICAL  COMPOUNDS. 

18.  Mix  two  or  three  grams  of  powdered  roll  sulphur 
and  an  equal  weight  of  very  fine  iron  filings  in  a  small 
mortar.    Examine  a  little  of  the  mixture  with  a  strong  lens 
or  microscope.     Can  you  distinguish  the  particles  of  iron 
and  of  sulphur?     Pass  a  small  magnet  over  the  mixture. 
Are   particles  of  iron  drawn  out   of  the  mixture  ?     Has 
chemical  action  taken  place  ?     Put  some  of  the  mixture  in 
a  dry  test  tube,  and  heat  it  with  the  flame  of  the  burner. 
Heat  gently  at   first  and   notice   the  changes.     At   first 
the  sulphur  melts  and  becomes  dark  colored,  but  soon  the 
whole  mass  begins  to  glow,  and  if  you  at  once  take  the 
tube  out  of  the  flame,  the  mass  will  continue  to  glow, 
becoming  brighter.     Now  allow  it  to  cool  down  and  then 
break  the  tube  and  put  the  contents  in  a  mortar.     Does 
the  mass  look  like  the  mixture  of  sulphur  and  iron  with 
which  you  started  ?     Examine  with  the  lens  and  with  the 
magnet.     Is  the  substance  now  a  mixture  or  a  chemical 
compound  ?     Why  ? 

19.  Mix  together  in  a  mortar,  as  intimately  as  possible, 
three   grams   of   zinc   dust   and    i^   grams   of  flowers  of 
sulphur.     Examine  the  mixture  with  a  lens  of  sufficient 
power  to  show  the  grains  of  sulphur  lying  side  by  side  with 
the  grains  of  zinc.     To  a  portion  in  a  test  tube  add  dilute 
hydrochloric  acid.     Notice  what  happens.     Make  with  the 
rest  of  the  mixture  a  conical  pile  on  a  piece  of  iron  or  bit 
of  asbestos  paper.     Apply  the  flame  of  a  match,  being 
careful  not  to  stand  too  near  the  mixture.     What  evidence 
is  there  that  chemical  change  takes  place  ?    Examine  with 
the  lens  the  powder  which  is  left,  to  see  whether  there  is 
zinc  or  sulphur  in  it.     Test  some  of  it  with  hydrochloric 
acid.     How  does  it  differ  from  the  original  mixture  ? 


HOMOGENEOUS   MATTER.  9 

20.  Examine  a  piece  of  granite.  Is  it  a  mechanical 
mixture  or  a  compound  ? 

Write  in  your  notebook  the  definitions  for  mechanical 
mixtures  and  for  chemical  compounds. 


HOMOGENEOUS  MATTER. 

Are  heterogeneous  bodies,  such  as  mechanical  mixtures, 
regarded  as  distinct  substances  in  chemistry?  What  do 
we  mean  by  the  term  homogeneous  substance  ? 

Make  a  list  of  homogeneous  and  heterogeneous  sub- 
stances in  your  notebook. 

Read  in  your  lecture  notes  and  text-books  in  regard  to 
the  meaning  of  the  terms  element,  compound,  and  physical 
mixture.  Give  examples  of  each  class. 

Would  it  be  possible  from  the  appearance  of  a  homo- 
geneous substance  to  tell  to  which  of  these  classes  it 
belonged  ?  How  would  you  determine  whether  a  given 
homogeneous  substance  was  an  element,  a  compound,  or  a 
physical  mixture  ? 

State  approximately  how  many  elements  are  known  at 
present. 

Read  in  your  text-book  about  the  relative  abundance  of 
the  elements. 

Which  is  the  most  abundant  element  ? 

Which  elements  make  up  three  fourths  of  the  solid  crust 
of  the  earth  ? 

Which  elements  are  the  chief  constituents  of  plants  and 
animals? 

Look  over  the  list  of  elements  on  page  1 14,  and  notice 
how  the  symbols  have  been  formed. 

In  what  respects  do  physical  mixtures  differ  from 
mechanical  mixtures  ? 

REISER'S  LAB.  CHEM. — a 


10  LABORATORY   WORK   IN   CHEMISTRY. 

THE  CHEMICAL  BALANCE. 

21.  Examine  the  construction  of  the  analytical  balance. 
Find  out  from  the  instructor  how  to  use  it.     Weigh  accu- 
rately a  clean  watch  crystal.     The  object  to  be  weighed 
must  always  be  placed  on  the  left-hand  scale  pan.     Never 
handle  the  weights  with  the  ringers,  but  always  use  the 
pincette  provided  for  the  purpose.     In  placing  weights  on 
the  right-hand  scale  pan  follow  the  directions  of  the  in- 
structor, who  will  also  explain  to  you  how  to  adjust  the 
rider  and  to  count  up  the  weights  properly. 

22.  Determination  of  Specific  Gravity.  —  Select  a  small 
piece   of   some    solid   which  is  insoluble  in   water,  —  for 
example,  a  piece  of   marble  or   a   bit    of  iron   or   brass. 
Attach  a  short  piece  of  thread  to  it  so  that  it  can  after- 
wards be  suspended  in  water  contained  in  a  small  beaker. 
Weigh  it  carefully  in  the  air.     Then  suspend  it,  by  means 
of  the  thread,  from  the  hook  above  the  scale  pan  so  that  it 
is  submerged  in  water  in  a  small  beaker.     Determine  its 
weight  in  the  water.     Calculate  the  specific  gravity. 

Determine  also  the  specific  gravity  of  a  liquid,  such  as 
alcohol,  using  a  small  specific  gravity  bottle.  First  weigh 
the  clean,  dry  bottle,  then  weigh  it  filled  .with  alcohol. 
Wash  out  the  alcohol  with  water,  and  now  weigh  it  filled 
with  water.  Calculate  the  specific  gravity. 

MEASUREMENT  OF  LIQUIDS. 

23.  For  measuring  liquids  the  liter  and  its  subdivisions 
are  used.     Find  out  the  relation  between  the  unit  of  weight 
and  the  unit  of  capacity  in  the  metric  system.     How  is  the 
liter  related  to  the  meter  ? 


THE  AIR.  II 

Compare  the  liter  flask  with  the  half-  and  quarter-liter 
flasks.  Inquire  of  the  instructor  how  to  use  the  pipettes, 
burettes,  and  graduated  cylinders.  Measure  the  capacity 
of  two  large  glass  bottles.  In  measuring  liquids  be  care- 
ful that  your  eye  is  in  the  same  plane  with  the  surface. 

24.  Construct  a  graduated  test  tube  by  running  in  accu- 
rately 5  or  10  cubic  centimeters  of  water  from  a  burette. 
Mark  the  height  at  which  the  water  stands  with  a  strip  of 
adhesive  paper.  Then  add  another  10  cubic  centimeters, 
mark  as  before,  and  so  continue  until  within  half  an  inch 
of  the  top.  Afterwards  make  permanent  marks  on  the 
glass  with  a  file  or  a  diamond. 


THE  AIR. 

[LABORATORY  DEMONSTRATION.] 

25.  Air  has  Weight.  —  To  determine  roughly  the  approxi- 
mate weight  of  air,  select  two  flasks  of  about  250  cubic 
centimeters'  capacity  which  have  been  blown  in  the  same 
mold  and  are  therefore  of  equal  size.  Fit  them  tightly 
with  corks.  Cork  one  permanently  and  reserve  it  as  a 
counterpoise.  Add  to  the  other  about  25  cubic  centi- 
meters of  water,  and  boil  the  water  over  a  lamp  until  the 
interior  is  full  of  steam  and  all  air  has  been  expelled ;  then 
quickly  cork  tightly,  remove  the  lamp,  and  allow  the  flask 
to  cool.  Tare  now  the  second  flask  with  the  first  and  such 
additional  weight  as  may  be  necessary.  [In  weighing  the 
flasks,  use  a  large  lecture  table  balance.]  Draw  the  cork 
of  the  flask  from  which  the  air  has  been  expelled,  and  after 
the  air  has  entered,  determine  the  increase  in  weight. 
Determine  the  volume  of  the  air  weighed  by  adding  water 


12  LABORATORY   WORK   IN   CHEMISTRY. 

from  a  graduated  measure  until  the  flask  is  filled  to  the 
former  level  of  the  cork.  From  these  values  the  weight  of 
a  liter  of  air  under  the  conditions  of  the  experiment  can  be 
calculated.  Read  in  a  text-book  on  physics  how  the  accu- 
rate value  for  the  weight  of  a  liter  of  air  has  been  determined. 

26.  The  pressure  of   the  atmosphere  is  measured   by 
means  of  the  barometer.     Describe  in  your  notebook  how 
a  barometer  is  made.     Read  the  barometer  in  the  labora- 
tory, and  enter  the  value  in  your  notebook,  together  with 
the  date  and  time  of  day. 

27.  In  a  small  clay  or  porcelain  crucible  put  a  small 
piece  of  zinc.     Support  the  crucible  on  a  pipe-stem  triangle, 
and  heat  it  with  the  burner.     After  the  zinc  is  melted,  stir 
it  with  an  iron  wire.     Continue  to  heat  and  stir  until  it  is 
no  longer  liquid.     Describe  what  has  taken  place. 

Now  repeat  the  experiment,  but  before  heating  the  zinc, 
add  enough  dry  borax  powder  to  form  a  complete  cover 
over  the  metal  when  both  are  melted.  It  may  be  neces- 
sary to  heat  the  crucible  with  the  blast  lamp.  Is  the  zinc 
changed  to  a  powder  in  this  case  ?  What  conclusion  can 
you  draw  from  this  experiment  ? 

28.  Metals  when  Heated  or  when   Burnt   in   the  Air 
Increase  in  Weight.  —  Place  a  saltspoonful  of   pulverized 
iron  in  a  small  crucible,  then  put  in  a  short  piece  of  iron 
wire  that  will  serve  as  a  stirring  rod,  and  now  weigh  the 
crucible  with  the  prescription  scales.     Note  the  weight  in 
your  laboratory  book.     Heat  the  crucible  with  the  burner, 
stir  the  powder  with  the  wire,  taking  care  not  to  lose  any 
of  the  powder.     Reweigh  when  cold,  and  describe  in  your 
notes  what  has  happened. 


THE  AIR. 

29.    Repeat  the  preceding  experiment,  using  about 
a  gram  of  magnesium  powder  instead  of  the  iron. 


of 


[LABORATORY  DEMONSTRATION.] 

30.  A  Candle  and  Other  Combustibles  in  Burning  Increase 
in  Weight.  —  On  one  pan  of  the  prescription  scales  place 
a  short  candle,  and  directly  above  it  suspend  a  wide  glass 
tube  or  lamp  chimney  containing  pieces  of  caustic  soda,  a 
substance  that  has  the  power  of  absorbing  the  products  of 
combustion  of  the  candle.  See  Fig.  2.  Bring  the  appara 


FIG.  2. 

tus  into  equilibrium  with  weights  ;  then  light  the  candle 
and  let  it  burn  for  a  few  minutes.  When  the  apparatus 
has  become  cold,  it  will  be  found  to  be  heavier  than  before 
the  candle  was  burnt.  What  inference  can  be  drawn  from 
this  and  the  two  preceding  experiments  ? 

31.  Burning  Substances  only  Combine  with  a  Portion  of 
the  Air.  —  Float  an  evaporating  dish  on  water  in  the  pneu- 
matic trough.  Put  a  piece  of  phosphorus  a  little  larger 
than  a  pea  in  the  dish.  Observe  great  care  in  handling  the 
phosphorus,  using  the  forceps  and  cutting  it  under  the 
water.  Set  fire  to  the  phosphorus  with  a  hot  wire,  and 
quickly  place  a  glass-stoppered  bell  jar  ovejjtjie  burning 

^R^ 

0*  TBM 


14  LABORATORY    WORK   IN   CHEMISTRY. 

phosphorus,  and  let  the  bell  jar  dip  an  inch  or  two  into 
the  water.  At  first  some  air  will  be  driven  out  of  the  bell 
jar  on  account  of  the  expansion  due  to  heat.  After  the 

burning  has  stopped,  allow  it 
to  stand  for  some  time.  The 
fumes  gradually  disappear, 
and  an  invisible  gas  remains. 
Notice  that  the  water  has  risen 
in  the  bell  jar.  Bring  the  bell 
jar  into  such  a  position  that 
the  height  of  the  water  is  the 
same  on  the  outside  and  in- 
side. Take  out  the  stopper 
and  insert  a  burning  splinter 

or  a  burning  candle.  Does  it  continue  to  burn  ?  What 
conclusions  do  you  draw  from  the  results  of  this  experi- 
ment in  regard  to  the  composition  of  the  air  ? 

MEASUREMENT  OF  GASES. 

32.  Gases  are  measured  in  eudiometers  and  in  gas  meas- 
uring tubes.  Notice  how  these  are  constructed.  Ther- 
mometers, a  barometer,  and  pneumatic  troughs  are  also 
required. 

To  compare  one  gas  volume  with  another,  it  is  necessary 
that  they  should  be  at  the  same  temperature  and  under  the 
same  pressure.  Why  is  this  so  ? 

What  effect  does  increase  of  temperature  have  upon  the 
volume  of  a  gas  ? 

What  is  the  law  of  Charles  ?    (Gay-Lussac's  Law  ?) 

What  effect  do  changes  of  pressure  have  upon  the  vol- 
ume of  a  gas  ? 

What  is  Boyle's  law  ?    (Mariotte's  Law  ?) 


MEASUREMENT  OF  GASES.  15 

The  standard  temperature  and  pressure  for  comparing 
gas  volumes  is  o°  Centigrade  and  760  millimeters  of 
mercury  pressure. 

33.  From  the  laws  of  Charles  and  Boyle,  deduce  the 
following  formula  for  reducing  gas  volumes  to  standard 
conditions :  — 

VP 
V'= — 

760(1+—} 

v     273; 

In  this  formula 

Vf=the  volume  of  the  gas  under  standard  conditions. 

V  =the  volume  of  the  gas  under  a  pressure  of  P  milli- 
meters of  mercury  and  at  a  temperature  of  /  degrees. 

34.  Measure  a  volume  of  air  in  a  eudiometer  over  mer- 
cury.    Read  the  temperature  by  means  of  a  thermometer 
that   dips   into    the   mercury.     Read   the   height   of   the 
barometer.     Notice  whether  the  mercury  is  at  the  same 
level  on  the  outside  and  inside  of  the  eudiometer.     How 
does  this  difference  of  level  affect  the  pressure  upon  the 
gas  ?     Measure  the  difference  of  level  and  determine  the 
pressure  of   the  gas.      Now   calculate    what   the   volume 
would  be  under  standard  conditions. 

The  usual  method  of  determining  the  weight  of  a  volume 
of  gas  is  to  calculate  what  the  volume  would  be  under 
standard  conditions,  and  then  to  multiply  by  the  weight  of 
one  volume  under  standard  conditions.  See  table  of  the 
weights  of  gases,  page  115. 

A  gas  that  is  saturated  with  moisture  occupies  a  larger 
volume  or  exerts  a  greater  pressure  than  it  would  if  it  were 
perfectly  dry.  It  is  really  a  mixture  of  the  dry  gas  with 
the  vapor  of  water.  To  reduce  such  a  gas  volume  to  stand- 
ard conditions,  it  is  necessary  to  subtract  from  the  apparent 


i6 


LABORATORY  WORK  IN  CHEMISTRY. 


pressure  of  the  gas  the  pressure  which  the  aqueous  vapor 
exerts  at  the  observed  temperature.  Then  proceed  as  in 
the  previous  case.  See  Table  of  Tension  of  Aqueous 
Vapor,  page  116. 

OXYGEN,  0. 

35.  Preparation  of  Oxygen  from  Mercuric  Oxide.  —  Place 
in  an  ignition  tube  a  layer  of  mercuric  oxide  about  ^  of  an 
inch  in  thickness.  Close  the  tube  with  a  stopper  provided 
with  a  delivery  tube.  Support  the  tube  with  the  clamp 


FIG.  4. 

stand  in  the  position  shown  in  Fig.  4,  so  that  the  delivery 
tube  is  under  the  cylinder  filled  with  water  in  the  pneumatic 
trough.  Heat  the  ignition  tube,  and  collect  the  gas  in  gas 
bottles.  What  is  formed  on  the  sides  of  the  ignition  tube  ? 
When  the  evolution  of  gas  has  nearly  stopped,  take  the 
delivery  tube  out  of  the  water,  then  take  away  the  flame 
from  the  ignition  tube.  Examine  the  gas  by  inserting 
a  stick  with  a  spark  on  the  end.  What  takes  place  ? 
Is  the  gas  contained  in  the  vessel  ordinary  air  ? 

36.  Repeat  the  preceding  experiment,  using  potassium 
chlorate  instead  of  mercuric  oxide. 

37.  Weigh  off  10  grams  of  potassium  chlorate,  and  grind 
it  in  the  mortar  to  a  fine  powder;  then  mix  with  it  inti- 
mately  10  grams  of  powdered   manganese  dioxide.     Fill 
the  ignition  tube  one  half  full  of  the  mixture,  heat,  and  fill 
all  your  gas  bottles  with  oxygen. 


OXYGEN.  17 

38.  Physical  Properties  of  Oxygen.  —  The  gas  in  the 

bottles  is  invisible.  The  slight  cloud  which  appears  when 
the  gas  is  first  collected,  is  due  to  the  presence  of  a  small 
quantity  of  solid  particles  that  have  been  carried  over  by 
the  gas.  After  standing  a  short  time  this  cloud  disappears. 
The  gas  is  tasteless  and  inodorous.  (Inhale  a  little  from 
one  of  the  bottles.)  It  is  slightly  heavier  than  the  air. 
When  subjected  to  a  sufficiently  low  temperature  and  high 
pressure  it  becomes  liquid. 

39.  Chemical  Properties  of  Oxygen.  —  Into  a  jar  of  oxy- 
gen introduce  a  little  lump  of  roll  sulphur  by  means  of  a 
deflagrating  spoon.     Let  it  stand  a  short  time,  and  notice 
what  change,  if  any,  takes  place.     Repeat  this  experiment, 
using  charcoal  and  phosphorus  in  place  of  the  sulphur. 
[Phosphorus   should  be  handled  with  great  care.      It  is 
always  kept  under  water.      If  a  small  piece  is  wanted,  take 
out  a  stick  with  the  forceps  and  put  it  under  water  in  an 
evaporating  dish  or  pneumatic  trough.      While  it  is  under 
water  cut  off  a  piece  of  the  size  wanted.      Take  this  out 
by  means  of  the  forceps,  lay  it  for  a  moment  on  a  piece  of 
filter  paper  to  absorb  the  water,  then  quickly  put  it  in  a 
deflagrating  spoon.]     What  does  this  experiment  show  ? 

40.  In  a  deflagrating  spoon  set  fire  to  a  little  sulphur 
and  let  it  burn  a  short  time  in  the  air. 

Notice  whether  it  burns  with  ease  or 
with  difficulty.  Notice  the  odor  of  the 
fumes.  Now  plunge  the  burning  sul- 
phur into  a  bottle  of  oxygen  (Fig.  5). 
Does  it  burn  more  readily  in  oxygen 
than  in  the  air?  Notice  the  odor  of 
the  fumes  in  the  bottle.  Is  it  the  same  FIG-  5- 

as  that  noticed  when  the  burning  took  place  in  the  air  ? 


1 8  LABORATORY   WORK   IN   CHEMISTRY. 

41.  Fasten  a  bit  of  charcoal  to  a  wire,  heat  it  in  the 
flame  until  it  is  red  hot,  then  plunge  it  into  oxygen.     How 
does  the  burning  in  oxygen  compare  with  the  burning  in 
air? 

42.  Burn  a  small  piece  of  phosphorus  in  the  air  and  in 
oxygen.     In  the  latter  case  the  light  emitted  from   the 
burning  phosphorus  is  so  intense  that  it  is  painful  to  some 
eyes  to  look  at  it.    After  the  burning  is  over  let  the  bottle 
stand.     Does  it  become  clear  ? 

43.  Insert  a  stick  with  a  spark  on  the  end  into  the 
bottles  in  which  sulphur,  charcoal,  and  phosphorus  have 
been   burnt.    Does   it   continue  to   burn  ?    Is   there   any 
oxygen  left  ?    What  do  you  infer  from  this  ? 

44.  Straighten  a  steel  watch  spring  by  taking  hold  of 
the  ends  and  slowly  passing  it  through  the  flame  until  all 
parts  have  been  red  hot.     Wind  a  little  thread  around  one 
end  and  dip  it  into  melted  sulphur.     Set  fire  to  this  and 
plunge  it  into  a  vessel  containing  oxygen.     For  a  moment 
the  sulphur  will  burn,  but  soon  the  steel  will  begin  to  burn 
brilliantly,  and  the  burning  will  continue  as  long  as  there 
is  oxygen  left  in  the  vessel.     The  phenomenon  is  of  great 
beauty,  especially  if  observed  in  a  dark  room.     The  walls 
of  the  vessel  become  covered  with  a  dark  reddish-brown 
substance.     Some  of  the  products  of  combustion  will  also 
be  found  at  the  bottom  in  the  form  of  globules. 

45.  Into  another  jar  of  oxygen  plunge  a  lighted  strip  of 
magnesium  ribbon.    Hold  the  magnesium  with  the  forceps, 
and  use  a  piece  about  four  inches  long. 


OXYGEN.  19 

46.  Fasten  a  short  piece  of  a  candle  to  a  bent  wire  so 
that  it  can  be  lowered  into  a  gas  bottle   (Fig.   6). 
Light  the  candle  and  insert  into  a  jar  of  oxygen. 
Notice  the  appearance  of  the  flame.     How  does  the 
burning  in  oxygen  compare  with  the  burning  in  air? 
Keep  the  jar  in  which  the  candle  is  burning  covered, 

and  notice  what  happens  after  a  time.     Why  is  the 
candle  extinguished  ?  FIG.  6. 

Write  in  your  notebook  a  full  account  of  all  the  experi- 
ments you  have  made  with  oxygen. 

What  inference  can  you  draw  from  these  experiments 
in  regard  to  the  nature  of  combustion  ? 

State  the  characteristic  physical  and  chemical  properties 
of  oxygen. 

Read  in  the  text-books  about  the  slow  action  of  oxygen 
at  ordinary  temperature.  Give  examples  of  substances  that 
are  acted  upon  by  oxygen  at  ordinary  temperature. 

[LABORATORY  DEMONSTRATION.] 

47.  Experiment  to  show  that  Oxygen  Disappears  when 
Phosphorus  Burns  in  it.  —  Float  a  small  dry  evaporating 
dish  on  the  water  of   the  pneumatic  trough.     Put  a  dry 
piece  of  phosphorus  as  large  as  a  bean  in  the  dish.     Place 
over  it  a  glass-stoppered  bell   jar  of  two  liters'  capacity. 
Remove  the  stopper  and  sink  the  bell  jar  so  far  into  the 
water  that  it  is  half  full  of  water  and  the  level  is  the  same 
outside  and  inside.     Fasten  the  bell  jar  securely  in  this 
position,  and  now  displace  the  air  in  it  with  oxygen.     This 
can  be  done  by  depressing  the  oxygen  delivery  tube  until 
it  is  under  the  edge  of   the  bell  jar  and  then  allowing  a 
slow  current  of  the  gas  to  enter.     After  about  two  liters 
of  oxygen  have  been  passed  in,  the  stopper  having  been 


20  LABORATORY   WORK   IN   CHEMISTRY. 

removed  during  this  time,  it  may  be  assumed  that  the 
air  has  all  been  displaced.  Now  heat  the  end  of  a  piece 
of  wire  red  hot,  plunge  it  into  the  bell  jar  and  set  fire 
to  the  phosphorus  in  the  dish.  Quickly  withdraw  the 
wire  and  instantly  close  the  bell  jar  tightly  with  the 
stopper.  At  first  the  heat  expands  the  gas,  but  it  soon 
begins  to  contract,  and  finally,  if  enough  phosphorus  was 
present,  and  all  the  air  was  displaced,  the  water  will 
completely  fill  the  vessel. 

What  has  become  of  the  oxygen  ?  What  would  have 
happened  if  under  the  same  conditions  sulphur  and  char- 
coal had  been  burnt  in  oxygen  ?  If  iron  or  magnesium 
had  been  used  in  place  of  phosphorus  ?  If  the  phosphorus 
had  been  burnt  in  air  contained  in  the  bell  jar  instead  of 
in  oxygen  ?  In  which  previous  experiment  have  you  tried 
this,  and  what  was  the  result?  From  all  the  experiments 
with  oxygen,  what  do  you  conclude  in  regard  to  the  nature 
of  combustions  in  air  and  in  oxygen  ? 

THE  CONSERVATION  OF  MATTER. 

48.  Select  two  small  beakers  and  fill  one  of  them  about 
one  third  full  of  a  solution  of  barium  chloride ;  fill  the 
other  one  third  full  of  dilute  sulphuric  acid.  Place  both 
beakers  in  the  scale  pan  of  a  prescription  balance,  and 
bring  it  into  equilibrium.  Now  carefully  pour  one  solution 
into  the  other ;  do  not  lose  a  drop,  and  place  the  empty 
beaker  upon  the  scale  pan.  Has  a  chemical  change  taken 
place  ?  Has  there  been  any  change  in  weight  ? 

What  is  the  law  of  the  conservation  of  matter  ? 

Can  energy  be  created  or  destroyed  ?  Give  reasons  for 
your  answer.  State  the  law  of  the  conservation  of  energy. 
What  does  it  mean  ? 


DEFINITE  PROPORTIONS. 


21 


LAWS  OF  CONSERVATION  OF  MATTER  AND  DEFINITE 
PROPORTIONS. 

49.    Analysis  of  a  Known  Weight  of  Potassium  Chlorate. 

—  Select  a  clean  piece  of  hard  glass  tubing  about  five 
inches  long  and  not  more  than  \  of  an  inch  in  outside  diam- 
eter. Draw  out  one  end  in  the  flame  and  seal  it.  Round 
the  edges  of  the  other  end  in  the  flame.  When  cold,  weigh 
the  empty  tube  with  the  analytical  balance.  Note  the 
weight  in  your  laboratory  notebook.  Put  into  the  tube  a 
few  small  crystals  of  potassium  chlorate.  [Less  than  -£$ 
of  a  gram  should  be  taken.]  Weigh  again.  The  differ- 
ence between  the  second  and  first  weighings  is  the  amount 
of  chlorate  in  the  tube.  Push  into  the  middle  of  the  tube 
a  small  plug  of  glass  wool  or  asbestos  fiber.  Weigh  again, 
and  then  bend  the  tube  as  shown  in  the  figure.  When 
the  tube  is  cold,  connect  it  by  means  of 
a  capillary  tube  and  two  rubber  tubes 
with  the  gas  measuring  burette  as  shown 
in  Fig.  7.  The  rubber  tube  that  con- 
nects with  the  burette  should  be  pro- 
vided with  a  pinchcock,  and  the  burette 
should  be  full  of  water  to  the  o  mark. 
Now  open  the  pinchcock  and  notice 
whether  the  joints  are  air  tight.  If  they 
are,  the  water  in  the  burette  will  not  fall, 
although  the  water  in  the  pressure  tube 
stands  but  an  inch  or  two  above  the 
base.  If  the  joints  are  air  tight,  heat  the 
crystals  gently  with  a  small  flame,  and 
continue  heating  until  no  more  gas  is 
evolved.  Let  the  apparatus  stand  for 
half  an  hour,  or  until  it  has  acquired  the  FIG.  7. 


22  LABORATORY   WORK   IN   CHEMISTRY. 

temperature  of  the  room.  Read  the  height  of  the  barome- 
ter ;  take  the  temperature  of  the  gas  by  means  of  a  ther- 
mometer attached  to  the  outside  of  the  tube.  Bring  the 
level  of  the  water  in  the  burette  and  the  pressure  tube  to 
the  same  height  and  read  the  volume  of  gas.  From  the 
observed  volume,  temperature,  and  pressure  calculate  the 
weight  of  oxygen  that  has  been  obtained  from  the  known 
weight  of  potassium  chlorate.  Disconnect  the  tube  con- 
taining the  potassium  chloride  and  weigh  it.  Calculate 
the  weight  of  potassium  chloride  that  is  left.  The  weight 
of  oxygen  found  added  to  the  weight  of  potassium  chloride 
should  equal  the  weight  of  potassium  chlorate  taken. 

KC103  =  KC1  +  03. 

A  repetition  of  the  experiment  with  a  different  quantity  of 
potassium  chlorate  will  show  that  this  compound  has  a 
definite  and  invariable  composition. 

What  is  the  law  of  definite  proportions  ?  Find  out  from 
your  instructor  and  your  text-books,  how,  from  the  laws 
of  conservation  of  matter  and  definite  proportions,  chemi- 
cal changes  can  be  represented  by  means  of  equations  like 
the  one  above. 

NITROGEN,  N. 

50.  Nitrogen  is  most  conveniently  prepared  by  burning 
phosphorus  in  air  confined  in  a  bell  jar  over  water.  The 
phosphorus  forms  a  white,  solid  oxide  which  readily  dis- 
solves in  the  water.  Prepare  nitrogen  in  this  way,  using 
the  apparatus  described  in  Experiment  31,  under  The  Air. 
What  is  the  color,  odor,  and  taste  of  the  gas  ?  Try  the 
effect  of  introducing  into  it,  successively,  several  burning 
bodies,  as,  for  example,  a  candle,  a  piece  of  sulphur,  phos- 
phorus, etc.  Do  these  substances  continue  to  burn  ? 


Of 


WATER.  23 

r+'m> 


VOLUMETRIC  COMPOSITION  OF 


51.  A  gas  burette,  such  as  was  used  in  the  analysis  of 
potassium  chlorate,  and  a  phosphorus  absorption  pipette 
are  required.     One  hundred  cubic  centimeters  of  air  under 
atmospheric  pressure  and  temperature  are  taken  into  the 
burette  and  transferred  to  the  absorption  pipette.      The 
phosphorus  oxidizes  slowly,  and  the  air  is  filled  with  white 
fumes.     After  some  minutes  these  disappear.     Now  draw 
back  the  residual  nitrogen  and  measure  the  volume.    Again 
pass  the  gas  into  the  pipette,  let  it  remain  a  short  time, 
and  then  bring  it  back  into  the  burette.     If  no  further  con- 
traction has  taken  place,  all  the  oxygen  has  been  removed. 

Repeat  the  analysis  until  good  results  are  obtained. 
Is  the  air  a  chemical  compound,  or  a  physical  mixture? 
What  are  your  reasons  for  regarding  it  as  such  ? 

WATER,  H.,0. 

52.  Occurrence  and   Distribution.  —  Read  in  your  text- 
book about  the  occurrence  of   water.       Many  apparently 
dry  substances  contain  water.     Put  into  a  dry  test  tube  a 
small  piece  of  wood,  and  heat  gently.     What  evidence  do 
you  obtain  that  water  is  given  off?     Do  the  same  thing 
with  a  piece  of  fresh  meat. 

53.  Water  of  Crystallization.  —  Heat  gently  a  few  crys- 
tals of  alum  in  a  test  tube.     Describe  what  takes  place. 

Perform  a  similar  experiment  with  some  gypsum,  which 
is  the  natural  substance  from  which  plaster  of  Paris  is 
made. 

Heat  also  in  a  tube  a  few  small  crystals  of  copper  sul- 
phate, or  blue  vitriol.  In  this  case  the  loss  of  water  is 


24  LABORATORY    WORK   IN   CHEMISTRY. 

accompanied  by  the  loss  of  color.  After  all  the  water  is 
driven  off,  the  powder  left  behind  is  white.  On  dissolving 
it  in  water,  however,  the  solution  will  be  seen  to  be  blue, 
and  if  the  solution  is  evaporated,  blue  crystals  will  be 
again  obtained. 

54.  Efflorescent  Substances.  —  Select  a  few  clear,  trans- 
parent crystals  of  sodium  sulphate,  or  Glauber's  salt.     Put 
them  in  a  watch  glass,  and  let  them  lie  exposed  to  the 
air  for  an  hour  or  two.     They  soon  lose  their  luster,  and 
become  white  and  powdery  on  the  surface.     What  is  the 
cause  of  this  ? 

55.  Deliquescent  Substances.  —  Expose  a  few  pieces  of 
calcium  chloride   to  the  air.      Its  surface  will  soon  give 
evidence  of  being  moist,  and  after  a  time  the  substance 
will  dissolve  in  the  water  which  is  absorbed. 

56.  Physical  Properties  of  Water.     Melting  and  Boiling 
Points.  —  For  this  experiment  a  thermometer  with  a  scale 
from  —5°  to  above  100°  is  required.     Place  some  broken 
ice  in  a  beaker  and  insert  the  thermometer ;  let  it  remain 
for  some  time,  then  read  the  temperature.    Does  the  quan- 
tity of  ice  melting  make  any  difference  in  the  temperature 
as  indicated  by  the  thermometer  ? 

Boil  water  in  a  beaker,  and  take  the  temperature  with 
the  thermometer.  Does  the  quantity  of  water  that  is  boiled 
or  the  rapidity  of  boiling  make  any  difference  in  the  tem- 
perature ? 

How  does  the  temperature  of  the  ice  compare  with  that 
of  the  lambent  water,  or  the  temperature  of  boiling  water 
with  that  of  the  steam  above  it  ?  What  becomes  of  the 
heat  that  enters  the  containing  vessel  ? 

What  is  the  effect  of  a  change  of  atmospheric  pressure 


WATER.  25 

upon  the  boiling  point  ?     What  effect  do  changes  of  pres- 
sure have  upon  the  melting  point  ? 
What  is  the  specific  gravity  of  ice  ? 

57.  Distillation  of  Water.  —  Use  for  the  purpose  a  glass 
retort     holding    about 

200  cubic  centimeters. 
Dissolve  a  few  crys- 
tals of  copper  sul-  '  "lIlC^WiL.  '  f^V 
phate  in  100  cubic 
centimeters  of  water. 
Pour  the  solution  into 
the  retort.  Fasten  the 
latter  in  the  retort 
stand.  Over  the  neck 
of  the  retort  push  the 
neck  of  a  glass  flask,  and  allow  the  flask  to  dip  in  cold 
water  contained  in  the  pneumatic  trough.  See  Fig.  8. 
Distil]  the  solution  until  only  a  little  is  left  in  the  retort. 
What  becomes  of  the  copper  sulphate  ?  What  would  hap- 
pen if  a  volatile  substance  had  been  dissolved  in  the  water 
instead  of  the  non-volatile  copper  sulphate  ? 

[LABORATORY  DEMONSTRATION.] 

58.  Decomposition  of  Water  by  Means  of  an  Electric 
Current.  —  Connect  the  wires  from  the  poles  of  a  Bunsen 
battery  with  the  platinum  electrodes  of  a  U-tube  apparatus 
for  decomposing  water.     The  apparatus  is  filled  with  water, 
to  which  has  been  added  one  tenth  its  volume  of  sulphuric 
acid.     Start  the  current,  and  gas  bubbles  will  collect  in 
each  of  the  two  branches  of  the  tube.     After  the  current 
has  been  going  for  a  few  minutes,  notice  that  there  is  twice 
as  much  gas  in  one  branch  as  in  the  other.     When  one 

KEISER'S  LAB.  CHEM. — 3 


26  LABORATORY   WORK   IN   CHEMISTRY. 

branch  is  full  of  gas,  stop  the  current.  Hold  a  lighted 
match  above  the  stopcock  of  the  branch  containing  the 
larger  gas  volume.  Open  the  stopcock  and  light  the 
escaping  gas.  A  flame  will  be  noticed.  The  gas  burns. 
Is  it  air,  or  oxygen,  or  nitrogen?  Bring  a  splinter  with  a 
glowing  end  over  the  other  stopcock  and  allow  the  gas 
to  escape.  Does  the  gas  act  like  oxygen  ?  Describe  this 
experiment  in  your  notebook. 

HYDROGEN,  H. 

59.  Preparation  of  Hydrogen.  —  Throw  a  small  piece  of 
sodium  *  on  water.  While  it  is  floating  on  the  surface  apply 
a  lighted  match  to  it.  What  takes  place  ?  The  flame  is 
due  to  burning  hydrogen.  The  yellow  color  of  the  flame 
is  caused  by  the  presence  of  sodium,  some  of  which  also 
burns.  If  a  piece  of  filter  paper  is  floated  on  the  water, 
and  the  sodium  is  then  thrown  upon  the  moist  paper,  the 
hydrogen  will  take  fire  spontaneously.  This  is  also  the 
case  if  the  sodium  is  thrown  upon  warm  water. 

A  small  piece  of  potassium  thrown  upon  water  decom- 
poses it  in  the  same  way ;  the  hydrogen  takes  fire  even 
when  the  metal  is  thrown  upon  cold  water,  and  burns  with 
a  violet  flame.  The  color  is  due  to  the  presence  of  potas- 
sium, some  of  which  also  burns.  Fill  a  test  tube  with 
water  in  the  pneumatic  trough,  and,  by  means  of  a  sharp- 
pointed  file,  pick  up  a  small  piece  of  sodium  and  quickly 
plunge  it  under  the  opening  of  the  tube.  The  sodium 
will  ascend  to  the  top,  and  in  this  way  the  gas  can  be 

1  Sodium  and  potassium  must  be  handled  with  care.  It  is  not  advisable  to 
use  a  piece  larger  than  a  grain  of  wheat.  Sodium  and  potassium  are  preserved 
under  oil.  Phosphorus  is  kept  under  water.  Do  not  make  the  mistake  of  try- 
ing to  cut  sodium  or  potassium  under  water. 


HYDROGEN.  2? 

collected.  Add  more  sodium  until  the  tube  is  full  of  gas. 
Put  your  thumb  over  the  mouth  of  the  tube,  bring  it  out 
of  the  water,  and  test  the  gas  with  a  lighted  match.  Does 
it  burn  ? 

Test  the  water  in  the  trough  with  red  litmus  paper. 
What  change  do  you  notice? 

Explain  the  chemical  changes  that  have  taken  place. 

[LABORATORY  DEMONSTRATION.] 

60.  Certain  metals  which  do  not  appreciably  decompose 
water  at  ordinary  temperature,  or  which  act  upon  it  slowly, 
decompose  it  easily  at  elevated  temperatures.  This  is  true 
of  iron  and  zinc.  If  steam  is  passed  through  a  tube  con- 


FIG.  9. 


taining  iron  turnings  or  fine,  bright  iron  wire  nails  heated 
to  redness,  the  water  is  decomposed,  the  oxygen  is  retained 
by  the  iron  in  chemical  combination,  while  hydrogen  is 
liberated.  Zinc  dust  decomposes  water  at  100°.  Fill  a 
hard  glass  combustion  tube  with  zinc  dust,  and  pass  a 
current  of  steam  through  the  tube.  Gently  heat  the 
tube  in  a  furnace,  and  collect  the  hydrogen  over  the 


28  LABORATORY   WORK   IN   CHEMISTRY. 

water  in  the  pneumatic  trough.     See  Fig.  9.     Light  the 
hydrogen  and  observe  the  color  of  the  flame. 

Magnesium  powder,  when  heated  in  a  current  of  steam, 
burns  brilliantly ;  hydrogen  is  given  off  and  may  be  col- 
lected as  in  the  case  of  the  zinc. 

61.  Hydrogen  prepared  from  Acids.  —  Into  a  test  tube, 
put  a  few  pieces  of  granulated  zinc,  and  pour  upon  the  zinc 
enough  hydrochloric  acid  to  cover  it.  What  do  you  notice  ? 
After  the  action  has  continued  for  a  minute  or  two,  apply 
a  lighted  match  to  the  mouth  of  the  tube.  What  takes 
place  ?  Try  the  same  experiment  with  dilute  sulphuric 
acid.  What  is  the  result  ?  The  gas  given  off  is  hydrogen. 
For  the  purpose  of  .collecting  the  gas  the  operation  is  best 
performed  in  a  flask  of  about  250  cubic  centimeters'  capacity. 
Fit  a  rubber  stopper  with  two  holes  into  the  neck  of  the 

flask.  Into  one  of  these 
holes  put  a  funnel  tube, 
which  must  reach  nearly 
to  the  bottom  of  the  flask. 
Into  the  other  put  a  glass 
tube  bent  at  right  angles. 
This  tube  need  only  pass 
through  the  stopper  and 
not  extend  below  it.  The 
outside  end  of  this  tube  is 

connected  by  means  of  a 
FIG.  10.  J 

short  piece  of  rubber  tub- 
ing, with  a  delivery  tube  which  dips  under  the  water  in 
the  pneumatic  trough.  [Instead  of  the  bottle  shown  in  the 
figure  an  ordinary  flask  may  be  used,  provided  care  is 
taken  not  to  break  it  in  putting  in  the  zinc.]  Put  a  small 
handful  of  granulated  zinc  into  the  flask,  and  pour  upon  it 


HYDROGEN.  29 

enough  of  a  cooled  mixture  of  sulphuric  acid  and  water 
(one  volume  of  concentrated  acid  poured  into  six  volumes 
of  water  —  never  pour  water  into  concentrated  sulphuric 
acid  !}  to  cover  it.  Usually  a  brisk  evolution  of  gas  will 
take  place  at  once.  Wait  two  or  three  minutes,  and  then 
collect  some  of  the  gas  by  displacement  of  water.  Should 
the  action  become  slow,  add  a  little  more  dilute  acid.  Fill 
several  cylinders  and  bottles  with  the  gas. 

Explain  the  chemical  change  that  takes  place  in  the 
flask.  Evaporate  some  of  the  liquid  in  the  flask,  and 
obtain  zinc  sulphate. 

62.  Physical    Properties    of    Hydrogen.  —  Notice    that 
hydrogen  is  a  colorless,  tasteless,  inodorous  gas.     When 
made  by  the  action  of  zinc  on  acids,  it  has  a  slightly  dis- 
agreeable odor.      This  is  due  to  the  presence  of   small 
quantities  of  impurities.     If  these  are  removed,  the  odor 
disappears.     This  can  be  done  by  passing  the  gas  through 
a  solution  of  potassium  permanganate. 

Hydrogen  is  not  poisonous  and  may  be  inhaled  with 
impunity.  It  is  the  lightest  substance  known.  The  air  is 
14.4  times  heavier ;  oxygen  is  16  times  heavier.  Place 
a  vessel  containing  hydrogen  with  the  mouth  upward 
and  uncovered.  In  a  short  time  examine  the  gas  in 
the  vessel  and  see  whether  it  is  hydrogen.  What  has 
happened  ? 

63.  Hydrogen  Poured  Upward.  —  Hold  a  glass  cylinder, 
containing  air,  inverted  in  one  hand ;  with  the  other  hand 
bring  a  cylinder  full   of    hydrogen  from   the   pneumatic 
trough,    and   gradually   pour   the    hydrogen    up   into   the 
inverted  cylinder.    The  air  will  be  displaced.    On  examina- 
tion the  inverted  cylinder  will  be  found  to  contain  hydro- 
gen, while  the  one  with  the  mouth  upward  will  contain 


LABORATORY   WORK  IN  CHEMISTRY. 


none.  Soap  bubbles  or  small  collodion  balloons  filled  with 
hydrogen  rise  in  the  air.  Fill  a  collodion  balloon  with 
hydrogen,  using  the  Kipp's  apparatus  to  generate  the  gas. 
When  full,  close  the  opening  by  means  of  a  string,  and  let 
it  go.  Large  balloons  are  always  filled  with  hydrogen  or 
some  other  light  gas.  Some  kinds  of  illuminating  gas  are 
rich  in  hydrogen,  and  may  be  used  for  the  purpose. 

What  is  the  weight  of  a  liter  of  hydrogen  under  the 
standard  conditions  ?  What  is  the  weight  of  a  liter  of 
oxygen  under  the  same  conditions  ?  Divide  the  latter  by 
the  former. 

[LABORATORY  DEMONSTRATION.] 

64,  Diffusive  Power  of  Hydrogen.  —  A  porous  clay  bat- 
tery jar  is  closed  with  a  cork,  and  a  long  glass  tube  open 

at  both  ends  is  passed  through 
the  cork.  Fasten  the  apparatus 
in  a  stand,  and  make  the  end  of 
the  glass  tube  dip  into  water  con- 
tained in  a  beaker,  or  arrange  it 
as  shown  in  Fig.  11.  Now  care- 
fully bring  a  bell  jar  full  of  hydro- 
gen over  the  battery  cup,  and 
notice  what  happens.  Take  away 
the  hydrogen  vessel,  and  observe 
what  takes  place.  Write  out  a 
description  and  explanation  of 
this  experiment  in  your  note- 
book. 

What  is  the  law  of  the  diffu- 
sion of  gases  ? 

What  are  the  relative  rates  of 
FJG.  ii.  diffusion  of  hydrogen  and  oxygen  ? 


HYDROGEN.  31 

65.  Chemical    Properties  of   Hydrogen.  —  Hold  a  wide- 
mouthed  bottle  or  cylinder  filled  with  hydrogen  with  the 
mouth  downward.     Insert    into  the  vessel  a 

long  lighted  taper.  The  gas  takes  fire  at  the 
mouth  of  the  vessel,  but  the  taper  is  ex- 
tinguished. On  withdrawing  the  taper  and 
holding  the  wick  for  a  moment  in  the  burning 
hydrogen,  it  will  take  fire  ;  but  on  putting  it 
back  in  the  hydrogen,  it  will  be  again  ex- 
tinguished. What  does  this  experiment  show? 
Other  burning  bodies  behave  in  the  same  way. 

66.  Pass  hydrogen  from  a  generating  flask,  or  Kipp's 
apparatus,  through  a  (J-tube  containing  granular  calcium 
chloride,  to  dry  the  gas.     Connect  by  means  of   rubber 
tubing  the  nozzle  of  a  blowpipe  with   the  [J-tube.     After 
the   hydrogen    has    been   allowed    to   escape   for    a   few 
moments,  set  fire  to  it.     [Always  be  cautious  in  working 
with  hydrogen.     The  danger  consists  in  the  fact  that  a 
mixture  of  hydrogen  and  oxygen,  or  hydrogen  and  air,  is 
explosive.     It  requires  a  flame   or   spark   to  explode   it. 
Always  let  the  gas  escape  for  a  time,  and  collect  a  test 
tube  full,  and  light  it  to  see  if  it  will  burn  quietly  before 
applying    a   flame    to   the   hydrogen    escaping   from   the 
apparatus.]     In  a  short  time  it  will  be  seen  that  the  flame 
is  practically  colorless,  and  gives  no  light.     That  it  is  hot 
is  shown  by  holding  a  piece  of  platinum  or  a  piece  of  some 
other  metal  in  it. 

Hold  a  dry  vessel  over  the  flame  of  burning  hydrogen. 
Drops  of  water  will  condense  on  its  surface  and  run  down. 
If  a  cold,  dry  bell  jar  is  held  above  the  flame,  the  steam 
formed  will  be  condensed  and  drops  of  water  will  collect  on 
the  walls  of  the  vessel. 


32  LABORATORY   WORK   IN  CHEMISTRY. 

What  conclusion  do  you  draw  from  this  experiment  in 
regard  to  the  composition  of  water  ? 

67.  A  strong  glass  cylinder  is  filled  one  third  full    of 
oxygen  in  the  pneumatic  trough.     The  rest  of  the  water 
is  then  displaced  by  hydrogen.     Take  it  out  of  the  water 
covered  with  a  glass  plate.     Remove  the  plate,  and  at  the 
same  instant  apply  a  light.     A  sharp  explosion  will  occur. 
What  is  the  cause  of  the  explosion  ? 

68.  Great  heat  is  evolved  when  hydrogen  and  oxygen 
combine.     A  jet  of  oxygen  issuing  into  a  flame  of  burning 
hydrogen  produces  an  intense  heat,  and  is  used  practically 
for  melting  platinum  and  for   producing   the   lime   light. 
The  apparatus  used  is  called  the  oxyhydrogen  blowpipe. 
Examine  the  construction  of  the  oxyhydrogen  blowpipe. 

[LABORATORY  DEMONSTRATION.] 

69.  Hold  in  the  flame  of  the  oxyhydrogen  blowpipe  suc- 
cessively a  piece  of  iron  wire,  a  steel  watch  spring,  a  piece 
of  copper  wire,  a  piece  of  zinc,  a  piece  of  platinum  wire. 

Allow  the  flame  to  strike  upon  a  piece  of  quicklime. 
Describe  the  effects  produced. 

Make  a  tabular  statement  of  the  characteristic  properties, 
both  physical  and  chemical,  of  hydrogen. 

COMPOSITION   OF   WATER   BY  VOLUME   AND    BY 
WEIGHT. 

[LABORATORY  DEMONSTRATION.] 

70.  Into  the  eudiometer  filled  with  mercury  and  stand- 
ing in  the  mercury  dish  as  shown  in  Fig.  13,  bring  about 
10  cubic  centimeters  of   oxygen.     Press   the   eudiometer 


COMPOSITION  OF  WATER. 


33 


down  on  the  conical  stopper  and  raise  the  pressure  tube 
until  the  level  of  the  mercury  is  the  same  outside  and 
inside  of  the  eudiometer.  Read  the  vol- 
ume of  oxygen  taken.  Bring  the  eudio- 
meter into  its  original  position  and  now 
introduce  not  less  than  25  cubic  centi- 
meters of  hydrogen  free  from  air.  Again 
read  the  volume  of  the  gas  when  it  is 
under  atmospheric  pressure.  Now  ex- 
pand the  gas  mixture  very  considerably 
by  lowering  the  pressure  tube,  and  close 
the  rubber  tube  with  a  strong  pinchcock. 
Explode  the  mixture  by  passing  an  elec- 
tric spark  through  the  tube.  Read  the 
volume  of  gas  that  remains,  and  calcu- 
late what  volume  of  hydrogen  has  united 
with  the  oxygen.  Test  the  residual 
gas  ;  it  will  be  found  to  be  hydrogen. 

Exactly  two  volumes  of  hydrogen 
unite  with  one  volume  of  oxygen.  If  either  gas  is  present 
in  excess  of  this  proportion,  the  excess  will  remain  uncom- 
bined. 

How  many  volumes  of  steam  are  formed  by  the  union 
of  two  volumes  of  hydrogen  with  one  of  oxygen  at  100°  ? 
How  could  this  be  shown  by  experiment  ?  Knowing  the 
composition  of  water  by  volume  and  the  relative  densities 
of  hydrogen  and  oxygen,  what  is  the  composition  of  water 
by  weight  ? 

In  Experiment  70,  water  may  be  used  instead  of  mercury 
to  fill  the  eudiometer  and  the  glass  dish  and  pressure  tube. 

Another  apparatus  that  may  be  used  for  this  experiment 
consists  of  a  gas  measuring  burette  and  an  explosion 
pipette,  as  shown  in  Fig.  14.  The  burette  is  filled  with 


FIG.  13. 


34 


LABORATORY   WORK  IN  CHEMISTRY. 


water;  then  about  10  cubic  centimeters  of  oxygen,  free 
from  air,  are  introduced  from  a  gasometer.  The  volume 
of  oxygen  is  read  off  when  it  is  under  atmospheric  pres- 


FlG.  14. 

sure.  About  25  or  30  cubic  centimeters  of  hydrogen  are 
then  introduced  and  the  volume  is  again  determined  under 
atmospheric  pressure.  [In  introducing  the  hydrogen  and 
oxygen  into  the  gas  burette  it  is  advisable  to  use  a  T-tube 
in  connecting  the  burette  with  the  gasometer  or  generat- 
ing flask.  By  this  means  it  is  possible  to  drive  out  all  air 
from  the  connections  before  admitting  the  gas  to  the 
burette.]  The  difference  between  the  first  and  second 
readings  gives  the  volume  of  hydrogen  taken.  Now  trans- 


COMPOSITION  OF  WATER. 


35 


fer  the  mixed  gases  to  the  explosion  pipette,  and  after 
closing  the  valves  explode  the  mixture  by  passing  a  cur- 
rent of  electricity  through  the  wires.  Draw  back  the 
residual  gas  into  the  burette  and  measure  its  volume. 
Test  it  by  means  of  a  lighted  match  on  expelling  it  into 
the  air,  to  show  that  it  is  hydrogen. 

71.  Gravimetric  Composition  of  Water. — Construct  an  ap- 
paratus like  that  shown  in  Fig.  15.  A  and  B  are  cylinders 
half  filled  with  concentrated  sulphuric  acid,  C  is  a  tube  of 


FIG.  15. 

* 

hard  glass,  the  middle  portion  of  which  is  filled  with  granu- 
lar copper  oxide.  The  copper  oxide  is  held  in  place  by 
plugs  of  glass  wool  or  asbestos  fiber.  D  and  E  are  tubes 
filled  with  calcium  chloride.  C  and  D  must  be  accurately 
weighed  before  the  experiment  on  the  analytical  balance. 
A  Kipp's  hydrogen  generating  apparatus  is  now  connected 
with  A  and  a  slow  current  of  hydrogen  is  allowed  to  drive 
out  all  the  air  in  the  apparatus.  When  this  has  been  done, 
heat  the  copper  oxide  in  C,  gently  at  first,  afterwards  more 
strongly.  Notice  that  the  copper  oxide  glows,  and  drops 
of  water  condense  in  D.  Keep  the  end  of  C  near  to  D 
warm,  from  the  beginning  of  the  experiment  until  the  end, 
so  that  water  does  not  condense  in  it.  When  all  the  cop- 
per oxide  is  reduced,  take  away  the  flame,  and  allow  the 
tube  to  cool  in  the  current  of  hydrogen.  Disconnect  C  and 


36  LABORATORY   WORK  IN   CHEMISTRY. 

D.  Draw  dry  air  through  them  to  displace  the  hydrogen 
and  weigh  them.  The  increase  in  the  weight  of  D  is  the 
weight  of  water  formed.  The  loss  in  weight  of  C  is  the 
weight  of  oxygen  contained  in  the  water. 

If  the  experiment  is  carefully  performed,  it  will  be  found 
that  the  weight  of  the  oxygen  is  eight  ninths  the  weight  of 
the  water. 

The  tube  D  can  be  used  several  times  without  refilling. 
When  not  in  use  it  is  closed  with  short  pieces  of  rubber 
tubing  plugged  with  glass  rods.  E  is  to  prevent  moisture 
of  the  air  from  entering  D,  and  A  and  B  serve  to  dry  the 
hydrogen  before  it  enters  the  tube  C.  It  is  advisable  to 
have  the  end  of  the  tube  C  touching  the  end  of  D  inside 
of  the  rubber  connection,  in  order  that  water  may  not  be 
condensed  on  the  rubber.  The  rubber  need  not  be  more 
than  J  of  an  inch  in  length. 

Read  in  one  of  the  larger  works  on  chemistry  an  account 
of  the  method  used  by  Dumas  in  determining  the  gravi- 
metric composition  of  water. 

What  is  meant  by  "  oxidation  "  and  "  reduction  "  ?  Give 
illustrations  of  each  kind  of  action. 


OZONE,  03. 

72.  Place  a  few  pieces  of  phosphorus  in  a  cylinder  or 
gas  jar  and  add  enough  water  to  partially  cover  it.  Cover 
the  vessel  and  allow  it  to  stand  for  20  minutes.  What 
evidence  is  there  that  the  phosphorus  is  slowly  oxidizing  ? 
Notice  the  odor  of  the  air  in  the  flask.  Boil  water  in  a 
small  beaker  or  test  tube  and  add  a  very  small  quantity  of 
starch  powder.  After  the  starch  has  dissolved,  add  a  crys- 
tal of  potassium  iodide.  Moisten  a  strip  of  filter  paper 
with  this  solution  and  introduce  it  into  the  cylinder  con- 


HYDROGEN   DIOXIDE.  37 

taining  the  phosphorus.  What  change  do  you  notice  ? 
Does  the  air,  or  oxygen,  or  nitrogen  act  upon  potassium 
iodide  and  starch  ?  What  has  caused  the  chemical 
change  ? 

73.  Make  a  solution  of  manganese  sulphate  in  water, 
dip  strips  of  filter  paper  into  it,  and  insert  them  into  the 
ozone  vessel.  What  change  takes  place  ? 

By  what  other  methods  can  ozone  be  obtained  ? 

Describe  the  manner  in  which  ozone  can  be  converted 
into  ordinary  oxygen. 


HYDROGEN  DIOXIDE,  H202,   AND  THE  LAW  OF 
MULTIPLE  PROPORTIONS. 

74.  Pour  10  cubic  centimeters  of  dilute  sulphuric  acid 
into  20  cubic  centimeters  of  water,  and  add  gradually  10 
grams  of  barium  dioxide.  Stir  constantly  with  a  glass  rod 
while  adding  the  powder.  Filter  the  solution.  The  clear 
filtrate  contains  hydrogen  dioxide.  Add  a  few  drops  of  the 
solution  to  a  solution  of  potassium  iodide  and  starch.  What 
effect  does  it  have  ?  To  another  portion  of  the  hydrogen 
dioxide  solution  add  a  small  quantity  of  powdered  man- 
ganese dioxide.  Notice  that  gas  bubbles  escape.  What 
is  the  cause  of  this  ?  To  another  portion  add,  drop  by 
drop,  a  very  weak  solution  of  potassium  permanganate. 
Does  the  violet  color  disappear  ? 

The  composition  of  hydrogen  dioxide  has  been  found  to 
be  one  part  of  hydrogen  to  16  parts  of  oxygen  by  weight. 
How  does  this  compare  with  the  composition  of  water  by 
weight  ?  How  do  the  properties  of  hydrogen  dioxide  com- 
pare with  those  of  water  ? 

What  is  the  law  of  multiple  proportions  ? 


38  LABORATORY  WORK  IN  CHEMISTRY. 

[LABORATORY  DEMONSTRATION.] 

75.  Experiments  to  illustrate  the  Law  of  Multiple  Pro- 
portions  Clean  a  small  porcelain  crucible,  support  it  on 

a  tripod  by  means  of  a  pipe-stem  triangle.  Heat  it  for  a 
few  moments  with  the  blue  flame  of  the  burner,  then  while 
it  is  still  warm,  place  it  by  means  of  pincers  in  a  desiccator 
containing  pieces  of  calcium  chloride.  After  the  crucible 
has  stood  in  the  desiccator  for  15  or  20  minutes,  weigh  it 
accurately  with  the  analytical  balance.  Note  the  weight 
carefully  in  your  notebook.  Cover  the  bottom  of  the 
crucible  with  a  layer  of  dry  copper  oxide  J  of  an  inch  in 
thickness.  Weigh  the  crucible  again.  The  increase  in 
weight  is  the  weight  of  copper  oxide  in  the  crucible. 
Place  the  crucible  again  on  the  pipe-stem  triangle,  and 
cover  it  with  a  lid  which  has  a  small  hole  in  the  middle. 
Through  this  hole  a  porcelain  tube  bent  at  right  angles 
is  passed.  (Rose's  crucible.)  The  end  of  the  tube  is 
connected  with  a  Kipp's  apparatus  for  hydrogen,  and  a 
slow  current  of  this  gas  is  turned  on.  After  a  minute  or 
two,  when  all  the  air  has  been  expelled  from  the  crucible, 
the  copper  oxide  is  heated  in  the  atmosphere  of  hydrogen 
for  8  or  10  minutes.  The  crucible  is  then  allowed  to  cool 
down  in  the  stream  of  hydrogen.  When  cold,  it  is  again 
weighed.  The  difference  between  the  third  and  second 
weighings  gives  the  weight  of  oxygen  in  the  known  weight 
of  copper  oxide.  Calculate  what  weight  of  copper  is  com- 
bined with  eight  parts  by  weight  of  oxygen  in  copper 
oxide. 

Now  repeat  the  experiment,  using  red  cuprous  oxide 
instead  of  black  cupric  oxide,  and  from  the  results  ob- 
tained calculate  the  weight  of  copper  that  is  in  combina- 
tion with  eight  parts  of  oxygen  in  this  compound.  How 


AMMONIA. 


39 


do  the  proportions  of  copper  in  the  two  compounds  com- 
pare with  each  other? 

Read  in  the  text-books  about  other  instances  of  chemi- 
cal combination  which  illustrate  the  law  of  multiple  pro- 
portions. 

AMMONIA,  NH3. 

76.  To  a  little  ammonium  chloride  on  a  watch  glass  add 
a  few  drops  of  a  strong  solution  of  caustic  soda,  an'd  notice 
the  odor  of  the  gas  given  off.     Do  the  same  thing  with 
caustic   potash.      Mix  small  quantities    of   quicklime  and. 
ammonium  chloride  in  a  mortar,  and  notice  the  odor.    Has 
ammonium  chloride  this  odor  ?     Explain  the  changes  that 
have  taken  place. 

What  is  the  principal  source  of  ammonia  compounds  ? 

77.  Arrange   an   apparatus   as 
shown  in  Fig.   16.      In  the  flask 
put   an    intimate   mixture   of    50 
grams  of  quicklime  and  25  grams 
of  ammonium  chloride.    Allow  the 
inverted  funnel  to  dip  just  below 

the    surface    of   the 

water  in  the  beaker. 

TT        ^i      ^11  FIG.  16. 

Heat  the  flask  on  a 

sand  bath.     After  the  air  is  driven  out,  the 
gas  will  be  completely  absorbed  by  the  water. 
Disconnect  the  funnel  and  turn  the  delivery 
tube  upward.     Bring  an  inverted  gas  cylinder 
FIG.  17.        over  the  tube,  and  collect  the  escaping  gas  by 
displacing  the  air.     As  the  gas  is  lighter  than  air,  it  is 
necessary  to  have  the  gas  jar  inverted,  and  the  delivery 
tube  should  extend  to  the  bottom.     See  Fig.  17.     The  jar 

O         I  J 


40  LABORATORY  WORK  IN  CHEMISTRY. 

in  which  the  gas  is  collected  should  be  dry,  as  water 
absorbs  ammonia  very  readily.  Hence,  also,  it  cannot  be 
collected  over  water.  Into  the  gas  thus  collected  intro- 
duce a  burning  stick  or  taper.  Does  it  burn  ?  Does  it 
support  combustion  ?  Avoid  breathing  any  quantity  of 
the  gas.  After  enough  has  been  collected,  pass  the  gas 
into  the  water  as  long  as  it  is  given  off. 

78.  Try  the  action  of  ammonia  upon  red  litmus  paper 
and  upon  tumeric  paper.     What  effect  does  it  have  upon 
blue  litmus  paper  ? 

79.  Dip  a  glass  rod  into  strong  hydrochloric  acid,  and 
bring  it  in  contact  with  ammonia  gas.     What  is  formed  ? 
Try  the  same  experiment,  using  strong  nitric  acid  in  place 
of  the  hydrochloric  acid.     What  is  formed  in  this  case  ? 

Make  a  tabular  statement  in  your  notebook  of  the 
physical  and  chemical  properties  of  ammonia.  What  is 
its  composition  by  volume  ? 

Read  in  a  text-book  how  ammonia  is  used  in  making  ice. 

NITRIC  ACID,  HN03. 

80.  Preparation  of  Nitric  Acid.  —  In  a  glass-stoppered 
retort  of  150  to  250  cubic  centimeters'  capacity  put  25  grams 
of  sodium  nitrate  (Chile  saltpeter)  and  25  grams1  of  con- 
centrated sulphuric  acid.    Push  over  the  neck  of  the  retort 
a  glass  flask,  and  allow  the  latter  to  dip  in  cold  water  con- 
tained in  the  pneumatic  trough.     See  Fig.  8,  p.  25.    On 

1  To  weigh  liquids,  balance  an  empty  beaker  on  the  prescription  scales; 
add  the  required  number  of  weights  to  the  scale  pan  opposite  the  one  con- 
taining the  beaker.  Then  pour  the  liquid  slowly  into  the  beaker  until  the 
pan  begins  to  sink. 


NITRIC  ACID.  41 

heating  gently,  nitric  acid  will  distill  over  and  be  con- 
densed in  the  flask.  In  the  latter,  stage  of  the  operation 
the  vessel  becomes  filled  with  a  reddish-brown  gas.  The 
acid  which  is  collected  has  a  somewhat  yellowish  color. 

Explain  the  chemical  change  that  takes  place  in  the 
retort. 

81.  Properties  of  Nitric  Acid.  —  Pure  nitric  acid  is  a 
colorless  liquid.  The  colored  acid  obtained  in  the  pre- 
ceding experiment  can  be  made  colorless  by  heating  it 
gently  in  the  flask  and  shaking  it.  When  the  acid  is  boiled, 
it  undergoes  slight  decomposition  into  oxygen,  water,  and 
compounds  of  nitrogen  and  oxygen.  One  of  these  is 
colored,  and  it  is  this  which  is  noticed  in  the  last  experi- 
ment, and  whenever  strong  nitric  acid  is  boiled.  Exposed 
to  the  light  for  a  long  time,  a  similar  decomposition  takes 
place. 

In  a  small  flask  put  a  few  pieces  of  granulated  tin. 
Pour  on  it  just  enough  strong  nitric  acid  to  cover  it. 
Heat  gently  with  a  small  flame.  What  takes  place  ?  What 
is  the  appearance  of  the  substance  left  in  the  flask  ?  It  is 
a  compound  of  tin,  hydrogen,  and  oxygen.  Dissolve  a  few 
pieces  of  copper  foil  in  concentrated  nitric  acid  diluted 
with  about  half  its  volume  of  water.  The  operation  should 
be  carried  on  in  a  flask  which  has  been  put  in  the  hood. 
What  action  takes  place  ?  After  it  is  over,  what  is  the 
appearance  of  the  liquid  in  the  flask  ?  Pour  it  out  and 
evaporate  to  crystallization  (be  careful  not  to  evaporate 
to  dryness).  Compare  the  substance  thus  obtained  with 
copper  nitrate.  Heat  specimens  of  each.  Treat  small 
specimens  with  sulphuric  acid.  Do  the  substances  appear 
to  be  identical  ?  What  reasons  have  you  for  considering 
them  identical  ? 

REISER'S  LAB.  CHEM. — 4 


42  LABORATORY   WORK  IN   CHEMISTRY. 

[LABORATORY  DEMONSTRATION.] 

82.  Pour  fuming  nitric  acid  into  a  wide  test  tube  so 
that  it  is  about  one  fourth  full.     Heat  the  end  of  a  stick 
of  charcoal  of  about  the  size  of  a  lead  pencil  and,  holding 
the  other  end  with  the  forceps,  introduce  the  heated  end 
into  the  acid.      It  will  continue  to   burn  with  a  bright 
light  even  though  placed  below  the  surface  of  the  liquid. 
The  action  is  oxidation.     The  charcoal  in  this  case  takes 
oxygen  from  the  acid,  and  not  from  the  air.     Care  must  be 
taken  in  performing  this  experiment,  that  the  hot  charcoal 
does  not  touch  the  sides  of  the  tube ;  and  a  large  beaker 
should  be  placed  beneath  the  test  tube,  so  that  in  case 
the  tube  should   break,   the   acid  would   be  caught   and 
prevented  from  doing  harm. 

83.  Put  a  small  piece  of  white  wool  or  worsted  into  an 
evaporating  dish,  add  a  few  drops  of  concentrated  nitric 
acid,   and  warm  gently.      Observe  the  yellow  stain.      A 
white  feather  is  stained  in  the  same  way.     Nitric  acid 
does  not  stain  cotton  or  calico  yellow. 

84.  To  a  dilute  solution  of  indigo  in  water  add  nitric 
acid.     Stir  the  mixture  and  keep  adding  the  acid  from 
time  to  time  until  the  indigo  is  all  bleached.     Repeat  the 
experiment,  using  litmus  solution. 

85.  Action  of  Ammonia  upon  Nitric  Acid.  —  To  a  dilute 
solution    of   nitric   acid   in  water   add   ammonia   solution 
until    the    liquid    smells    slightly   of    ammonia,    or    until 
a  piece  of   litmus   paper   dipped   into   the  liquid  is  col- 
ored blue.     Now  evaporate  on  the  water  bath.     A  solid 
white  crystallized  mass  remains  behind.     This  is  ammo- 
nium nitrate,  formed  by  the  union  of  the  ammonia  with 


NITROUS   OXIDE.  43 

the  nitric  acid.     Preserve  the  crystals  for  future  experi- 
ments. 

86.  When  nitric   acid   acts  upon  "metals,   nitrates   are 
usually  formed ;  for  example,  nitric  acid  acting  upon  cop- 
per gives  copper  nitrate.     See  Experiment  81.     In  these 
compounds  the  hydrogen  of  the  acid  has  been   replaced 
by  metals.      What  becomes  of  the  hydrogen  ?     Like  the 
pure  acid  the  nitrates  are  good  oxidizing  agents. 

Melt  a  few  crystals  of  potassium  nitrate  (saltpeter)  in 
an  ignition  tube,  then  drop  in  a  bit  of  charcoal  and  heat 
until  deflagration  takes  place. 

Repeat  the  experiment,  using  a  small  piece  of  roll  sul- 
phur in  place  of  the  charcoal.  Explain  what  -takes  place. 

87.  General  Reaction  for  All  Nitrates In  a  test  tube 

put  a   small   crystal   of  any  nitrate.      Add  a   few  drops 
of  pure  concentrated   sulphuric  acid.     Cool  the  mixture, 
and  when  cold,  pour  slowly  down  the  side  of  the  tube  a 
solution  of  ferrous   sulphate.      A  brown  ring   is   formed 
where   the  ferrous  sulphate    comes   in  contact  with  the 
heavier  sulphuric  acid  solution  at  the  bottom.     Ask  the 
instructor  for  an  explanation  of  the  reaction.     The  experi- 
ment is  frequently  made  use  of  in  the  analysis  of  unknown 
substances. 

What  are  the  distinguishing  properties  of  nitric  acid  ? 
How  is  it  prepared  commercially  ? 

NITROUS  OXIDE,  N20. 

88.  In  a  glass  retort  put  10  or  15  grams  of  ammonium 
nitrate.      Fix  the  retort  in  the  clamp  stand  so  that  the 
opening  of  the  neck  dips  under  the  water  of  the  pneu- 


44  LABORATORY  WORK  IN  CHEMISTRY. 

matic   trough.      Heat   the   ammonium   nitrate   gently  to 
secure  a  regular  evolution  of  the  gas.     Fill  several  bottles 
or  jars  with  the  gas  as  in  the  experiments  under  oxygen. 
Explain  what  change  takes  place  in  the  retort. 

89.  Insert  into  a  jar  of  the  gas  a  piece  of  wood  with  a 
live  coal  on  the  end.     Try  also  a  burning  candle.     Into 
another  bottle  of  the  gas  plunge  a  bit  of  burning  phos- 
phorus.     Also   introduce   burning    sulphur   into   nitrous 
oxide.     What  happens  in  each  case  ? 

90.  Fill  a  cylinder  half  full  of  nitrous  oxide,  and  then 
add  hydrogen  until  full.     Apply  a  light  to  the  mixture : 
it  explodes.     What  is  the  reaction  ?     Which  of  the  gases 
that  you  have  studied  does  nitrous  oxide  resemble  most 
nearly  ?     What  is  the  effect  of  inhaling  nitrous  oxide  ? 

91.  To  distinguish  nitrous  oxide  from  oxygen,  keep  a 
bottle  half  filled  with  the  gas  standing  in  the  pneumatic 
trough,  and  use  it  for  Experiment  94,  under  Nitric  Oxide. 

NITRIC  OXIDE,  NO. 

92.  Preparation  from  Copper  and  Nitric  Acid.  —  Use  an 

apparatus  like  that  used  for  preparing  hydrogen.  In  the 
flask  put  a  few  pieces  of  copper  foil.  Cover  this  with 
water.  Now  add  slowly,  waiting  after  each  addition,  ordi- 
nary concentrated  nitric  acid.  When  enough  acid  has 
been  added,  gas  will  be  given  off.  If  too  much  acid  is 
added  at  once,  it  not  infrequently  happens  that  the  evolu- 
tion of  gas  takes  place  so  rapidly  that  the  liquid  is  forced 
out  of  the  flask  through  the  funnel  tube.  What  is  the 
color  of  the  gas  in  the  flask  at  first  ?  What  is  it  after  the 


NITRIC  OXIDE.  45 

action  has  been  continued  for  some  time  ?     Collect  over 
water  two  or  three  vessels  full. 

What  is  the  chemical  change  which  causes  the  formation 
of  nitric  oxide  ? 

93.  Turn  one  of  the  vessels  containing  nitric  oxide  so  that 
the  mouth  is  upward,  and  uncover  it.      What  takes  place  ? 
Explain  the  appearance  of  the  colored  gas  in  the  previous 
experiment,  and  the  fact  that  it  afterward  disappeared. 
Avoid  inhaling  the  brown  fumes  that  are  formed  when 
nitric  oxide  comes  in  contact  with  air. 

94.  Pass  nitric  oxide  from  the  generating  flask  into  the 
bottle  containing  nitrous  oxide  mentioned  under  Experi- 
ment 91.     Are  brown  fumes  formed?     If  the  bottle  had 
contained  oxygen,  what  would  have  been  the  result  ? 

95.  Into  one  of  the  bottles  containing  nitric  oxide  insert 
a  burning  candle.     Does  the  gas  burn  ?     Does  it  support 
combustion  ? 

96.  Into  a  cylinder  containing  nitric  oxide  pour  a  few 
drops  of  carbon  disulphide,1  shake  the  cylinder,  and  then 
apply  a  burning  match.     A  brilliant  flame  is  produced. 
Substances  that  evolve  a  great  deal  of  heat  in  burning  will 
continue  to  burn  in  nitric  oxide.     When  substances  burn 
in  nitrous  and  nitric  oxide,  what  becomes  of  the  nitrogen 
in  these  gases  ? 

97.  Pass  a  current  of  nitric  oxide,  as  it  issues  from  the 
generating  flask,  into  a  solution  of  ferrous  sulphate.     The 

1  Be  careful  to  keep  the  bottle  containing  carbon  disulphide  away  from  the 
neighborhood  of  flames.  Carbon  disulphide  vapor  mixed  with  air  explodes 
when  lighted. 


46  LABORATORY  WORK  IN  CHEMISTRY. 

gas  is  retained  in  the  solution,  and  the  latter  becomes 
brown,  then  black  in  color.  If  the  dark  solution  is  now 
boiled,  the  color  again  becomes  yellow  or  yellowish  green. 
In  which  previous  experiment  has  this  brown  solution  of 
nitric  oxide  in  ferrous  sulphate  been  obtained  ? 

In  your  notebook  make  a  list  of  the  oxides  and  acids 
of  nitrogen.  Give  the  gravimetric  composition  of  each. 
What  important  law  do  they  illustrate  ?  Read  in  the  text- 
books how  the  other  oxides  are  prepared. 

VOLUMETRIC  COMPOSITION  OF  NITROUS  AND  NITRIC 
OXIDES,  AND  THE  LAW  OF  VOLUMES. 

[LABORATORY  DEMONSTRATION.] 

98.  Place  in  a  flask  of  200  cubic  centimeters'  capacity 
20  grams  of  potassium  nitrate  and  80  grams  of  ferrous 
sulphate.  Add  enough  water  to  cover  the  solids.  Close 
the  flask  with  a  tightly  fitting  rubber  stopper  carrying  a 
dropping  funnel  and  a  delivery  tube.  Allow  concen- 
trated sulphuric  acid  to  enter  the  flask,  drop  by  drop,  and 
heat  gently.  A  steady  stream  of  pure  nitric  oxide  is 
evolved.  Connect  the  delivery  tube  of  the  flask  by  means 
of  rubber  connections  with  a  J-tube,  and  connect  one  of 
the  open  ends  of  the  f-tube  with  a  gas-measuring  burette 
full  of  water,  such  as  was  used  in  the  volumetric  analysis  of 
the  air.  Allow  the  nitric  oxide  to  escape  from  the  open  end 
of  the  f-tube  for  5  or  10  minutes,  until  all  the  air  has  been 
displaced  in  the  flask  and  in  the  connections.  Then  close 
the  open  end  of  the  T-tube  with  a  rubber  tube  compressed 
with  a  pinchcock,  and  admit  the  gas  into  the  burette. 

Having  measured  100  cubic  centimeters  of  the  gas  in 
the  burette,  connect  the  latter  with  a  simple  absorption 
pipette  filled  with  water,  using  instead  of  the  ordinary 


CHLORINE.  47 

capillary  tube  a  tube  of  hard  glass  -J-  of  an  inch  internal 
diameter,  and  filled  with  granules  of  metallic  copper.  (The 
tube  is  filled  with  granular  copper  oxide,  and  then  the 
oxide  is  reduced  by  heating  in  a  stream  of  hydrogen.) 

The  tube  containing  the  copper  is  heated  to  a  dull  red 
heat  by  means  of  a  broad-top  burner,  and  the  nitric  oxide  is 
transferred  slowly  to  the  pipette.  The  gas  is  thus  decom- 
posed into  nitrogen,  the  oxygen  uniting  with  the  copper, 
forming  copper  oxide.  Draw  back  the  gas  into  the  burette, 
and  measure  the  volume  of  nitrogen.  How  does  its  vol- 
ume compare  with  the  volume  of  nitric  oxide  taken  ? 

Repeat  the  experiment,  using  nitrous  oxide  instead  of 
nitric  oxide.  Since  nitrous  oxide  is  soluble  in  water,  have 
the  water  of  the  burette  saturated  with  this  gas  before 
making  the  measurements.  What  volume  of  nitrogen  is 
contained  in  one  volume  of  nitrous  oxide  ?  What  is  the 
law  of  combination  of  gaseous  volumes  ?  Give  a  number 
of  examples  to  illustrate  the  law. 

CHLORINE,  Cl. 

99.  Preparation  of  Chlorine.  — A  flask  of  250  cubic  cen- 
timeters' capacity  provided  with  a  rubber  stopper  having  a 
funnel  tube  and  a  delivery  tube  is  required.  The  flask 
used  for  making  hydrogen  and  nitric  oxide  may  be  used. 
Put  into  the  flask  about  100  grams  of  powdered  man- 
ganese dioxide,  pour  upon  it  enough  ordinary  concentrated 
hydrochloric  acid  to  completely  cover  it.  Shake  the  mix- 
ture thoroughly,  so  that  the  acid  touches  all  parts  of  the 
bottom  of  the  flask.  Heat  very  gently  with  a  small  flame, 
and  chlorine  will  be  given  off.  Collect  several  dry  cylin- 
ders or  bottles  full  of  chlorine,  by  letting  the  delivery  tube 
extend  to  the  bottom  of  the  collecting  vessel,  and  covering 


48  LABORATORY  WORK  IN  CHEMISTRY. 

the  mouth  of  the  vessel  with  a  piece  of  paper  or  a  glass 
plate.  You  can  see  when  the  vessel  is  full  by  the  color  of 
the  gas.  [The  experiments  with  chlorine  should  be  carried 
on  in  a  good  hood.  Do  not  inhale  the  gas.'] 

By  what  reaction  is  the  chlorine  obtained  in  this  experi- 
ment ? 

100.  Properties  of  Chlorine.  —  Notice  the  color  and  odor 
of  the  gas.     Why  is  it  not  advisable  to  collect  the  gas  by 
the  displacement  of  cold  water  in  the  pneumatic  trough  ? 
Find  out  from  your  books  how  much  heavier  the  gas  is 
than  air.      Into   one  of  the  vessels  containing   chlorine 
introduce  a  little  finely  powdered  antimony.     What  be- 
comes of  the  antimony?    What  causes  the  evolution  of 
light  and  heat  ? 

101.  Into  a  second  jar  of  chlorine  introduce  a  piece  of 
very  thin  copper  foil  or  a  piece  of  Dutch  leaf.     What  evi- 
dence is  there  of  a  transformation  of  energy?     A  com- 
pound of  copper  and  chlorine  is  formed. 

102.  Into  a  third  vessel  insert  a  piece  of  paper  with 
writing  on  it,  some  flowers  and  pieces  of  colored  calico. 
Most  of  the  colors  will  be  destroyed  if  the  substances  are 
moist. 

103.  Into  a  fourth  vessel  introduce  a  dry  piece  of  the 
same  calico  as  that  used  in   Experiment   102.     The  dry 
piece  is  not  bleached,  the  moist  piece  is.     Explain  this. 

104.  Lower  a  burning  candle  into   a   cylinder  full  of 
chlorine.     See  Fig.  18.     Why  does  it  burn  with  a  smoky 
flame  ?     Does  the  chlorine  burn  ? 


HYDROCHLORIC  ACID. 


49 


105.  Add  chlorine  water,  made   by  bubbling    chlorine 
gas  through   cold  water,  to  a  dilute  solution  of  indigo. 
What   effect   does   it   have? 

Try  its   action    upon  litmus 
solution. 

106.  Moisten    a    strip    of 
filter  paper  with  turpentine, 
and  plunge  it  into  a  bottle  of 
chlorine.     See  Fig.  19.    Tur- 
pentine  is   a    compound    of 
carbon  and  hydrogen.      Ex- 
plain    what     takes     place. 
Usually    so    much    heat    is 

produced   that   the  substance   takes  fire. 
Experiment  104. 

What  are  the  distinguishing  properties  of  chlorine  ? 

Read  in  one  of  the  text-books  an  account  of  the  techni- 
cal methods  of  making  chlorine,  and  of  the  applications  of 
chlorine  for  bleaching  and  disinfecting  purposes.  How 
can  chlorine  be  liquefied  ?  Describe  in  your  notebooks 
Faraday's  method  of  liquefying  chlorine.  Why  are  oxygen, 
nitrogen,  and  hydrogen  so  much  more  difficult  to  liquefy 
than  chlorine,  ammonia,  and  nitrous  oxide  ?  What  is  the 
critical  temperature  of  a  gas  ? 


FIG.  18. 


FlG.ig. 


Compare  with 


HYDROCHLORIC  ACID,  HC1. 

[LABORATORY  DEMONSTRATION.] 

107.  Light  a  jet  of  hydrogen  in  the  air,  and  carefully 
introduce  it  into  a  vessel  containing  chlorine.  See  Fig.  20. 
Does  it  continue  to  burn  ?  What  is  the  appearance  of  the 


$O  LABORATORY   WORK   IN   CHEMISTRY. 

flame  ?    What  evidence  have  you  that  a  product  is  formed  ? 

Test  the  gas  remaining  in  the  jar  with  blue  litmus  solution 

shaken  up  in  it,  and  compare  with 
the  action  of  chlorine  on  the  solu- 
tion. 

What  does  this  experiment 
show  ?  What  is  the  composition 
of  hydrochloric  acid  by  volume  ? 
What  law  does  this  illustrate  ? 

108.    Preparation  of  Hydrochloric 
Acid  from  Salt  and  Sulphuric  Acid. 
—  Arrange  an  apparatus  similar  to 
FIG.  20.  that  used  in  preparing  a  solution 

of  ammonia  in  water,  but  use  in 

place  of  the  one-hole  rubber  stopper  one  with  two  holes, 
and  push  a  funnel  tube  through  the  second  hole.  Place 
50  grams  of  common  salt  in  the  flask.  Measure  20  cubic 
centimeters  of  water  into  a  beaker,  and  pour  into  it  50 
cubic  centimeters  of  strong  sulphuric  acid.  When  this 
mixture  has  cooled  down  to  the  ordinary  temperature,  pour 
it  upon  the  salt  in  the  flask.  Now  heat  the  flask  gently, 
and  a  regular  current  of  gas  will  be  given  off.  Conduct 
it  into  water ;  when  the  air  has  all  been  driven  out  of  the 
flask,  the  gas  is  all  absorbed  by  the  water  under  the  funnel. 
The  solution  of  the  gas  in  water  is  heavier  than  the  water, 
and  as  it  is  formed  you  can  see  it  settling  down  in  the 
beaker  and  mixing  with  the  dilute  solution.  After  the  gas 
has  passed  into  the  water  for  about  10  minutes,  discon- 
nect the  funnel.  Notice  the  fumes.  These  become  denser 
by  blowing  the  breath  upon  them.  Apply  a  lighted  match 
to  the  end  of  the  tube.  Does  the  gas  burn  ?  Collect 
some  of  the  gas  in  one  or  two  dry  cylinders  by  downward 


HYDROCHLORIC  ACID.  51 


displacement  of  air,  as  in  the  case  of  chlorine.  Then  con- 
nect the  generating  flask  again  with  the  funnel,  and  let 
the  heating  continue  until  no  more  gas  comes  over.  There 
is  left  in  the  flask  hydrogen  sodium  sulphate.  Write  in 
your  notebook  the  reaction  which  has  taken  place. 

109.  Properties  of  Hydrochloric  Acid. — Notice  the  physi- 
cal properties  of  the  gas,  its  color,  odor,  etc.      Insert  a 
burning  stick  or  candle  into  the  gas.     Does  it  support 
combustion  ? 

110.  Open  another  cylinder  filled  with  hydrochloric  acid 
gas  under  water  in  the  pneumatic  trough.    If  the  gas  con- 
tained no  air,  the  water  will  rush  into  the  cylinder  as  it 
would  into  a  vacuum.     What  is  "the  reason  for  this  ? 

111.  Moisten  a  strip  of  filter  paper  with  ammonia  solu- 
tion, and  plunge  it  into  a  cylinder  of  hydrochloric  acid  gas. 
Explain  what  takes  place. 

112.  Try  the  action  of  the  solution  of  hydrochloric  acid 
that  you  have  prepared  upon  some  granulated  zinc  in  a 
test  tube.     Is  a  gas  given  off  ?     What  is  it  ?     Add  some 
to  a  little  manganese  dioxide,  and  heat  gently.     Is  chloi  me 
liberated?     Add   10  or  12  drops  of  water  to  two  or  three 
drops  of  the  solution  of   the  hydrochloric  acid  and  taste 
it.     How  would   you   describe   the   taste  ?     Dip   a   piece 
of  blue  litmus  paper   into  the    dilute   hydrochloric   acid. 
Litmus  is  a  vegetable  color  prepared  for  use  as   a   dye. 
Other  vegetable  colors  are  changed  by  hydrochloric  acid. 
The  color  is  restored  by  adding  ammonia  or  caustic  soda 
solution.     What  is    the  composition  of  hydrochloric  acid 
by  weight  and  by  volume  ?     . 

When  hydrochloric  acid  is  treated  with  metals,  hydrogen 


52  LABORATORY   WORK  IN   CHEMISTRY. 

is  evolved,  and  chlorides  of  the  metals  are  formed ;  thus, 
with  zinc,  hydrochloric  acid  gives  zinc  chloride  and  hydro- 
gen ;  with  iron,  iron  chloride  and  hydrogen,  etc.  Nearly 
all  the  metallic  chlorides  are  soluble  in  water.  Silver 
chloride,  mercurous  chloride,  and  lead  chloride  are  insoluble 
in  water. 

113.  Add  a  drop  of  silver  nitrate  to  a  solution  of  any 
chloride.     What  is  precipitated  ?     What  remains  in  solu- 
tion ? 

Add  a  drop  of  mercurous  nitrate  to  a  solution  of  any 
chloride.  Interpret  the  reaction. 

Add  several  drops  of  lead  nitrate  to  a  solution  of  a 
chloride.  Explain  the  chemical  change. 

How  is  hydrochloric  acid  obtained  technically  ? 

ACIDS,  BASES,  AND  SALTS. 

Acids  and  bases  have  the  power  to  destroy  the  charac- 
teristic properties  of  each  other.  They  neutralize  each 
other.  The  most  common  acids  are  sulphuric,  nitric, 
and  hydrochloric  acids.  Among  the  more  common  bases 
are  caustic  soda,  caustic  potash,  lime  water,  and  ammonia 
water.  Acids  turn  litmus  red,  bases  turn  litmus  blue. 

114.  Make   a   dilute  solution   of   hydrochloric   acid   by 
measuring  10  cubic  centimeters  of  dilute  hydrochloric  acid 
from  the  reagent  bottle,  and  diluting  it  with  50  cubic  cen- 
timeters of  water.     Weigh  off  about  5  grams  of  sodium 
hydroxide    (caustic   soda),   and   dissolve    it    in    100   cubic 
centimeters  of  water.     Measure  20  cubic  centimeters  of 
the  dilute  acid,  by  means  of  a  burette,  into  a  beaker,  add 
a  drop  of  litmus,  and  now  with  another  burette  graduated 
into    fractions   of   cubic   centimeters  run  in  the   caustic 


ACIDS,    BASES,   AND   SALTS. 


soda  solution  until  the  liquid  just  turns  from  red  to  blue 
(Fig.  21).  Note  the  quantity  of  caustic  soda  that  was 
required  to  neutralize  20  cubic  centimeters  of  the  acid. 
Now  take  10  cubic  centimeters  of  acid,  and  determine  the 
amount  of  sodium  hydroxide  solution  re- 
quired to  neutralize  it.  Repeat,  using  15 
cubic  centimeters  of  acid.  What  relation 
do  the  quantities  of  alkali  used  bear  to  one 
another?  Are  they  in  the  proportion  of 
20:  10  :  15  ? 

Make  a  dilute  solution  of  sulphuric  acid 
and  of  caustic  potash  as  described  above, 
and  determine  the  quantity  of  alkali  re- 
quired to  neutralize  20,  10,  and  15  cubic 
centimeters  of  the  acid. 

Does  it  always  require  a  definite  quan- 
tity of  alkali  to  neutralize  a  definite  quan- 
tity of  an  acid  solution  ? 


FIG.  21. 


115.  When  acids  neutralize  bases,  salts  and  water  are 
formed.  Dissolve  10  grams  of  caustic  soda  in  100  cubic 
centimeters  of  water.  Add  hydrochloric  acid  slowly,  ex- 
amining the  solution  from  time  to  time  with  a  piece  of 
blue  litmus  paper.  As  long  as  the  solution  is  alkaline 
the  paper  stays  blue ;  the  instant  it  passes  the  point  of 
neutralization  the  paper  turns  red.  When  this  point  is 
reached,  evaporate  the  solution  to  complete  dryness,  and 
see  what  is  left.  Taste  the  substance  ;  has  it  an  acid  taste? 
If  it  is  common  salt  or  sodium  chloride,  how  ought  it  to 
conduct  itself  when  treated  with  sulphuric  acid  ?  Does  it 
conduct  itself  in  this  way  ?  Is  the  substance  an  alkali 
or  an  acid  ?  Is  it  neutral  ?  Write  the  equation  represent- 
ing its  formation. 


54  LABORATORY  WORK  IN   CHEMISTRY. 

116.  Repeat  the  experiment,  using  nitric  acid  instead  of 
hydrochloric   acid.      Compare    the   product    with    sodium 
nitrate  by  treating  a   small  specimen   of   each  with  sul- 
phuric acid  in  test  tubes.    Are  the  substances  identical? 
Write   the   equations   representing   the   neutralization  of 
nitric  acid  by  caustic  soda ;  sulphuric  acid  by  caustic  soda; 
hydrochloric  acid  by  caustic  potash  ;  nitric  acid  by  caustic 
potash.    How  could  you  readily  determine  whether  a  given 
substance   is  an  acid,  a  base,  or  a  salt  ?     Are  all  salts 
neutral  to  litmus? 

CARBON,  C. 

Read  the  chapter  on  carbon  in  your  text-book. 

In  how  many  forms  does  the  element  occur  ? 

What  are  the  characteristic  properties  of  the  diamond, 
graphite,  and  amorphous  carbon  ?  How  is  charcoal  made  ? 
How  are  coke,  lampblack,  and  boneblack  or  animal  char- 
coal obtained  ?  What  is  coal  ?  Charcoal  can  be  made  in 
the  laboratory  by  heating  wood  covered  with  sand  in  a 
crucible  until  gases  are  no  longer  given  off. 

117.  Decolorizing   Power   of    Boneblack.  —  To  a  dilute 
solution  of  indigo  add  a  tablespoonful  of  animal  charcoal. 
Heat  the  solution  to   boiling,  shaking  frequently  ;   filter 
while  hot.     If  the  filtrate  is  not  colorless,  run  it  through 
the   filter   several    times.     The   animal   charcoal   is   very 
porous,  and  the  coloring  matter  is  retained  in  the  pores. 

118.  Chemical  Conduct   of  Carbon.  —  At  ordinary  tem- 
peratures carbon  is  not  an  active  element ;   it  does  not 
unite  with  hydrogen,  oxygen,  chlorine,  or  nitrogen.     At 
high    temperatures    it    readily    combines    with     oxygen. 
Fasten  a  bit  of  charcoal  on  a  wire,  heat  it  in  a  flame  until 
red  hot,  then  plunge  it  into  a  bottle  filled  with  oxygen. 


MARSH  GAS.  55 


When  it  stops  burning  take  it  out  and  pour  into  the  bottle 
a  clear  solution  of  limewater.  The  limewater  becomes 
milky.  Explain  what  has  taken  place.  What  is  formed 
when  coal  burns  in  a  stove  or  in  a  grate? 

119.  At  high  temperatures  carbon  will  take  oxygen  from 
oxides.     Mix  intimately  in  a  mortar  two  or  three  grams  of 
powdered  copper  oxide  and  a  gram  of  powdered  charcoal. 
Heat  the  mixture   in  an   ignition  tube   provided  with  a 
stopper  and  outlet  tube,  so  that  the  gas  which  escapes  can 
be  passed  into  a  clear  solution  of  limewater  contained  in  a 
test  tube.     Is  a  white  precipitate  formed  ?    What  evidence 
have  you  that  oxygen  has  been  extracted  from  the  copper 
oxide  ?     When  cold,  pour  the  black  mixture  in  the  tube 
into  cold  water  in  a  beaker ;  stir  and  let  it  stand  a  short 
time.     A  red  powder  will  collect  on  the  bottom  of  the 
beaker.     What   is  it  ?     Write  the  equation   representing 
the  action  of  copper  oxide  upon  charcoal  at  high  tempera- 
tures.    Make  a  list  in  your  notebook  of  metals  that  are 
obtained  from  their  oxides  by  the  reducing  action  of  carbon. 

120.  Repeat    the    preceding    experiment,  using  white 
arsenic,  As2O3,  in  place  of  the  copper  oxide.     Write  the 
equation  representing  the  reaction. 

The  element  arsenic  is  volatile,  and  is  hence  driven  out 
of  the  bottom  of  the  tube  and  deposited  on  the  sides  above 
the  mixture  in  the  form  of  a  mirror  with  a  metallic  luster. 

MARSH  GAS,  CH4. 

121.  Preparation  of  Marsh  Gas.  —  Put  a  large  spoonful 
of  sodium  acetate  in  an  evaporating  dish.     Heat  with  a 
very  small  flame,  and  stir  until  the  water  of  crystallization 
has  been  driven  off.     At  first  the  salt  melts,  but  as  the 


56  LABORATORY  WORK  IN   CHEMISTRY. 

water  is  driven  off,  it  again  becomes  solid.  Now  weigh  off 
5  grams  of  potassium  hydroxide,  7^  grams  of  quicklime, 
and  5  grams  of  dry  sodium  acetate.  Powder  the  caustic 
potash  and  the  quicklime  in  the  mortar,  then  add  the 
acetate,  and  mix  intimately.  Put  the  mixture  into  an 
ignition  tube,  close  the  tube  with  a  stopper  carrying  a 
delivery  tube  (the  apparatus  used  in  making  oxygen  will 
do),  and  support  the  tube  horizontally  in  a  clamp  stand. 
Heat  the  mixture  and  collect  the  marsh  gas  in  the  bottles 
as  in  the  case  of  oxygen.  Notice  that  the  gas  is  colorless, 
transparent,  and  inodorous.  It  is  only  slightly  soluble  in 
water.  Apply  a  lighted  match  to  some  of  the  gas  in  a 
cylinder.  Does  it  burn  ?  Does  it  give  light  in  burning  ? 
What  is  formed  when  marsh  gas  burns  ? 

Read  in  the  text-books  about  the  occurrence  and  the 
properties  of  marsh  gas  ? 

ETHYLENE,  C^. 

122.  Preparation  of  Ethylene.  —  Put  a  few  pieces  of  gran- 
ulated zinc  in  a  test  tube,  add  alcohol  and  a  drop  or  two 
of  ethylene  bromide.  Warm  gently.  When  the  gas  is 
coming  off  rapidly,  apply  a  burning  match  to  the  mouth  of 
the  tube.  Does  the  ethylene  burn  with  a  bright,  lumi- 
nous flame  ?  What  is  formed  when  the  ethylene  burns  ? 


*      ACETYLENE, 

123.  Preparation  of  Acetylene.  —  Into  a  test  tube  pour 
5  cubic  centimeters  of  water,  then  drop  into  the  water  two 
or  three  pieces  of  calcium  carbide  of  about  the  size  of  a 
bean.  A  rapid  evolution  of  gas  takes  place.  After  the 
air  has  been  driven  out  of  the  test  tube,  apply  a  light  to 
its  mouth.  Does  the  acetylene  burn  ?  Is  the  flame  lumi- 


CARBON   DIOXIDE.  57 

nous  or  non-luminous  ?  Acetylene  mixed  with  air  or  oxy- 
gen explodes  violently  when  a  light  is  applied.  Write 
the  equations  representing  the  action  of  water  upon  calcium 
carbide,  and  of  oxygen  upon  acetylene  when  it  burns  in 
air.  Inquire  of  the  instructor  how  calcium  carbide  is  made. 
What  weight  of  carbon  is  united  with  one  part  by 
weight  of  hydrogen  in  marsh  gas,  ethylene,  and  acetylene  ? 
What  law  does  this  illustrate  ? 

CARBON  DIOXIDE,   C02. 

124.  Carbon  Dioxide   is   formed  in  Breathing.  —  Blow 
through  a  glass  tube  that  dips  into  clear  limewater  con- 
tained in  a  test  tube.     What  evidence  have  you  that  your 
lungs  give  off  carbon  dioxide  ? 

125.  Preparation  from  Carbonates  and  Acids. — To  dif- 
ferent test  tubes  containing  a  little  sodium  carbon-       _ 
ate  add  dilute  hydrochloric,  sulphuric,  nitric,  and        ^ 
acetic  acids.     Is  a  gas  given  off  ?     Test  the  gas  in 

each  tube  by  bringing  into  it  a  wire  with  a  small 
loop  on  the  end  (Fig.  22)  that  has  been  dipped 
into  clear  limewater.  The  drop  on  the  end  of  the 
wire  will  get  milky  if  carbon  dioxide  is  present. 
Treat  marble  with  any  acid,  and  test  the  gas 
evolved  in  the  same  way.  Is  it  carbon  dioxide  ? 
Write  the  reactions  in  each  case. 

126.  Arrange  an  apparatus  like  that  used  in  pre- 
paring chlorine,  or  as  shown  in  Fig.  10.      Put  into      O 
the  flask  some  pieces  of  marble,  and  pour  ordinary  FlG<  22' 
hydrochloric  acid  upon  it.     Collect  the  gas  by  the  down- 
ward displacement  of  air,  as  in  the  case  of  chlorine.     Fill 
several  bottles  with  the  gas. 

KEISER'S  LAB.  CHEM. — 5  %*• 


58  LABORATORY   WORK   IN   CHEMISTRY. 

127.  Properties  of  Carbon  Dioxide.  —  Notice  whether  the 
gas  has  color  or  odor.     Insert  a  lighted  candle  or  burning 

stick  into  the  gas.     What  takes  place  ?     Is 
the   gas   combustible  ?      Does    it    support 
combustion  ?    Pour  the  gas  from  one  of  the 
jars  upon  the  flame  of  a  burning  candle ; 
proceed  as  if  pouring  water  from  the  jar. 
See  Fig.  23.     Pour  some  of  the  gas  from 
one  vessel  to  another  and  show  that  it  has 
been  transferred.     Balance  a  beaker  on  the 
FIG.  23.          prescription  scales  and  pour  carbon  dioxide 
into  it.     If  the  balance  is  sensitive,  the  pan  on  which  the 
beaker  is  standing  will  go  down.     What  do  these  experi- 
ments show? 

,  What  is  the  specific  gravity  of  carbon  dioxide  com- 
pared with  air  ?  How  many  times  heavier  than  hydrogen 
is  it  ?  Is  it  easily  liquefied  ?  (Critical  temperature  ?) 

128.  Carbon  Dioxide  acts  upon  Bases  and  forms  Car- 
bonates. —  Pass  a  current  of  carbon  dioxide  into  a  solution 
of  caustic  soda  until  it  will  absorb  no  more.     Add  acid  to 
some  of  the  solution  thus  obtained,  and  convince  yourself 
that  the  gas  given  off  is  carbon  dioxide.     Write  the  equa- 
tions representing  the  reactions  which  take  place  on  pass- 
ing carbon  dioxide  into  the  caustic  soda  and  on  adding 
acid  to  the  solution.     Pass  carbon  dioxide  into  about  200 
cubic  centimeters  of  clear  limewater.     Filter  off  the  white 
insoluble  precipitate.    Try  the  action  of  a  little  acid  on  it. 
What   evidence  have  you  that  it    is  calcium  carbonate  ? 
How  could  you  easily  distinguish  limewater  from  caustic 
soda?     What  are  the  characteristic  properties  of  carbon 
dioxide  ?     When  carbon  is  burnt  in  oxygen,  what  relation 
is  there  between  the  volume  of  carbon  dioxide  and  the 


CARBON  MONOXIDE. 


-  _C\         59 


volume  of  oxygen  from  which  it  has  been  formed  ?  Is  it 
larger  or  smaller  ? 

129.  Calcium  carbonate  dissolves  in  a  solution  of  carbon 
dioxide  in  water.     Pass  carbon  dioxide  into  about  50  cubic 
centimeters  of  clear  limewater  ;  notice  that  the  precipitate, 
which  at  first  is  very  voluminous,  gradually  goes  into  solu- 
tion if  you  keep  on  passing  in  the  gas.     Acid  calcium  car- 
bonate, H2Ca(CO3)2,  is  formed,  and  this  is  soluble.     This 
compound  is  decomposed  on  heating  the  clear  solution,  and 
calcium  carbonate  is  again  precipitated.     Try  it. 

What   causes    the    "temporary   hardness"   of    natural 
waters  ?     How  can  it  be  removed  ? 

CARBON  MONOXIDE,  CO. 

130.  Preparation  from  Oxalic  Acid  and  Sulphuric  Acid. 

—  Put  10  grams  of  oxalic  acid  and  50  to  60  grams  of  con- 
centrated sulphuric  acid  into  a  250  cubic  centimeter  flask. 
Close  the  flask  with  a  stopper  provided  with  a  funnel  tube 
and  a  delivery  tube.  Connect  the  delivery  tube  with  two 
gas  washing  cylinders  containing  caustic  soda  solution. 
Heat  the  contents  of  the  flask  gently.  Carbon  dioxide 
and  carbon  monoxide  are  evolved.  In  passing  through 
the  caustic  soda  the  former  gas  is  absorbed.  Collect  the 
carbon  monoxide  in  gas  bottles  over  water.  Avoid  inhal- 
ing the  carbon  monoxide,  as  it  is  poisonous.  Light  some 
of  the  gas  in  a  cylinder  ;  notice  the  characteristic  blue 
flame.  What  is  formed  when  it  burns  ?  By  what  other 
methods  can  carbon  monoxide  be  prepared  ? 

131.  Carbon  Monoxide  a  Reducing  Agent.  —  Pass  carbon 
monoxide  over  some  heated  granules  of  copper  oxide  con- 
tained in  a  hard  glass  tube.     Is  the  oxide  reduced  ?     How 


60  LABORATORY   WORK  IN   CHEMISTRY. 

do  you  know  ?     Is  carbon  dioxide  formed  ?     How  can  it 
be  shown  ? 

In  what  respects  does  carbon  monoxide  resemble  and 
how  does  it  differ  from  carbon  dioxide  ?  How  can  one 
be  distinguished  from  the  other  ?  In  what  respects  does 
carbon  monoxide  resemble  and  how  does  it  differ  from 
hydrogen  ? 


THE    KINDLING   POINT:  INFLUENCE    OF    TEMPERATURE 
ON  CHEMICAL  ACTION. 

132.    Combustible   substances    must   be    raised   to   the 
kindling  temperature  before  they  will  burn.    When  cooled 


FIG.  24.  FIG.  25. 

below  this  temperature,  they  are  extinguished.  Light  a 
Bunsen  burner.  Bring  down  upon  the  flame  a  piece  of  wire 
gauze.  There  is  no  flame  above  the  gauze.  See  Fig.  24. 
That  the  gas  passes  through  unburned  can  be  shown  by 
applying  a  light  just  above  the  outlet  of  the  burner  and 
above  the  gauze.  The  gas  takes  fire  and  burns.  By  sim- 
ply passing  through  the  wire  gauze  the  gas  is  cooled  be- 
low its  burning  temperature.  Turn  on  a  Bunsen  burner. 
Do  not  light  the  gas.  Hold  a  piece  of  wire  gauze  about 
an  inch  and  a  half  or  two  inches  above  the  outlet.  Apply 
a  lighted  match  above  the  gauze,  and  the  gas  will  burn 
above  the  gauze,  but  not  below  it.  See  Fig.  25. 


THE  BLOWPIPE.  6 1 

133.  Take   apart   a   miner's    safety  lamp  and   examine 
its  construction.     Light  it,  put  it  together,   and  allow  a 
current  of  coal  gas  from  a  rubber  tube  to  strike  against 
the  wire  gauze.     The  coal  gas  will  burn  inside  the  lamp, 
but  the  flame  will  not  pass  through. 

What  change  do  the  combustibles  coal,  wood,  sulphur, 
etc.,  undergo  in  the  air  at  ordinary  temperatures  ?  At 
elevated  temperatures  ?  What  influence  has  temperature 
upon  the  rapidity  of  chemical  changes  ? 

ILLUMINATING  GAS. 

Read  in  the  text-books  how  illuminating  gas  is  made. 
What  gases  are  present  in  coal  gas  ?  How  is  water  gas 
made,  and  of  what  does  it  consist  ?  Which  of  the  gases 
present  in  coal  gas  and  in  water  gas  burn  with  a  luminous, 
and  which  with  a  non-luminous,  flame  ? 

THE  BLOWPIPE. 

134.  The  flame  of  the  Bunsen  burner  consists  of  two 
parts,  —  an  inner  flame  of  incomplete  combustion  and  an 
outer  flame  of  complete  combustion.     In  the  inner  flame 
there  is  an  excess  of  combustible  substances ;  these  are 
at  a  high  temperature,  and  will  take  away  oxygen  from 
certain  oxides  if  they  are  brought  into  the  flame.     The 
inner  flame  is  called  the  reducing  flame.     In  the  outer 
flame  there  is  an  excess  of  oxygen,  and  in  this  flame  cer- 
tain substances  undergo  oxidation,  hence  it  is  called  the 
oxidizing    flame.      The    blast-lamp   and   blowpipe   flames 
are  similar  to  the  Bunsen    burner ;  each  consists  of  an 
oxidizing  and  a  reducing  flame. 

Practice  with  the  blowpipe,  and  learn  to  blow  without 


62  LABORATORY   WORK  IN   CHEMISTRY. 

interrupting  the  blast  on  account  of  breathing.  Put  a 
small  piece  of  metallic  lead  in  a  depression  on  a  small 
piece  of  charcoal  and  heat  it  in  the  oxidizing  flame.  In 
making  the  oxidizing  flame,  hold  the  nozzle  of  the  blow- 
pipe in  the  middle  of  the  gas  flame.  Notice  the  formation 
of  lead  oxide. 

Mix  together  in  a  mortar  a  little  tin  oxide  and  sodium 
carbonate,  put  some  of  the  mixture  into  a  depression  on  the 
charcoal,  and  heat  it  in  the  reducing  flame.  In  making 
the  reducing  flame,  hold  the  blowpipe  nozzle  just  outside 
the  gas  flame. 

On  a  loop  of  platinum  wire  make  a  borax  bead ;  bring 
into  it  a  small  quantity  of  any  manganese  compound. 
Heat  in  the  oxidizing  flame  until  violet  in  color.  Now 
heat  in  the  reducing  flame  until  colorless.  Repeat  this 
several  times. 


THE   EQUIVALENT  WEIGHTS  OF  THE  ELEMENTS   AND 
THE  LAW  OF  RECIPROCAL  PROPORTIONS. 

The  equivalent  of  an  element  is  that  weight  of  the 
element  which  combines  with,  or  replaces  in  compounds, 
one  part  by  weight  of  hydrogen. 

135.    Determination  of  the  Equivalent  of  Zinc Weigh 

accurately  with  the  analytical  balance  from  .1  to  .2  grams 
of  zinc.  Place  the  zinc  in  a  small  porcelain  crucible, 
covered  with  a  small  inverted  funnel.  Both  crucible  and 
funnel  are  put  into  a  large  beaker  nearly  full  of  water,  as 
shown  in  Fig.  26.  A  gas  measuring  tube  of  100  cubic 
centimeters'  capacity  is  filled  half  full  of  dilute  sulphuric 
acid,  the  remaining  half  of  the  tube  is  filled  with  water, 
and  then  the  tube  is  closed  with  the  thumb  and  inverted 


MOLECULAR  WEIGHTS.  63 

over  the  crucible  and  funnel.  Hydrogen  is  evolved  and 
collects  in  the  tube ;  allow  the  action  to  continue  until 
all  the  zinc  has  dissolved.  [A  drop  of  copper 
sulphate  solution  added  to  the  dilute  acid  has- 
tens the  action,  but  at  the  same  time  introduces 
a  slight  error.]  Transfer  the  gas-measuring 
tube  to  a  large  cylinder  of  water,  and  when  it 
has  acquired  a  constant  temperature,  read  the 
volume  under  atmospheric  pressure,  the  tem- 
perature, and  the  height  of  the  barometer. 
Calculate  the  weight  of  hydrogen  obtained,  and 
the  equivalent  of  zinc. 


136.    Determine  in  a  similar  way  the  equiva- 
lent of  magnesium  and  aluminium.     Use  hydro- 
chloric acid  to  dissolve  the  aluminium.     Calcu-     FlG  ^ 
late  the  equivalent  of  oxygen  from  the  results 
obtained  in  analyzing  water.     See  Determination  of   the 
Gravimetric  Composition  of  Water,  Experiment  71. 

Ask  the  instructor  to  explain  to  you  how  the  equivalents 
of  other  elements  could  be  determined.  Make  a  table  of 
equivalent  weights  in  your  notebook. 

Is  the  equivalent  of  an  element  necessarily  a  constant 
quantity  ?  What  is  the  law  of  reciprocal  proportions  ? 
Give  a  number  of  examples  to  illustrate  it. 

MOLECULAR  WEIGHTS. 

What  is  the  theoretical  explanation  given  to  the  laws  of 
definite,  multiple,  and  reciprocal  proportions  ? 

What  is  the  hypothesis  of  Avogadro  ? 

State  briefly  the  principle  employed  in  determining 
molecular  weights.  What  is  the  atomic  weight  of  an 


64  LABORATORY  WORK   IN   CHEMISTRY. 

element  ?     How  is  the  equivalent  of  an  element  related 
to  its  atomic  weight  ? 

Read   in   the   text-books  how  the   atomic  weights  are 
determined. 

BROMINE,  Br. 

137.  Mix  together  about  a  gram  of  potassium  bromide 
and  two  grams  of  manganese  dioxide.      Pour  upon  the 
mixture  in  a  good-sized  test  tube  sufficient  dilute  sulphuric 
acid  to  cover  it.     Heat  gently.     What  takes  place  ?     In- 
terpret the  reaction.     Compare  with   the  preparation  of 
chlorine. 

138.  Dip  a  piece  of  moistened   litmus  paper  into  the 
bromine  vapor.      Is  it  bleached?      To  a  dilute  solution 
of  indigo  add  bromine  water  until  the  blue  color  disap- 
pears.    How  does  the  rapidity  of  bleaching  compare  with 
that  of  chlorine  ? 

139.  To  a  solution  of  potassium  bromide  add  chlorine 
water.     What   is   formed  ?     Add  a  few  drops   of   carbon 
disulphide,  shake,  and  observe  the  color  of  the  disulphide. 
Explain  what  has  taken  place. 

In  what  compounds  does  bromine  occur  in  nature  ? 
Read  the  chapter  on  bromine  in  your  text-books,  and 
make  a  list  of  the  distinguishing  properties  of  the  element. 

HYDROBROMIC  ACID,  HBr. 

140.  Preparation  of  Hydrobromic  Acid.  —  In  a  test  tube 
put  a  few  crystals  of  potassium  bromide.    Pour  on  them  a 
few  drops   of   concentrated   sulphuric   acid.      Notice   the 
white  fumes  of  hydrobromic  acid ;  at  the  same  time  some 


PREPARATION  OF  AMMONIUM   BROMIDE.  6$ 

reddish  brown  vapor  of  bromine  is  also  given  off.  Tre'at 
a  few  crystals  of  potassium  or  sodium  chloride  in  the  same 
way.  What  difference  is  there  between  the  two  cases? 
When  hydrobromic  acid  and  sulphuric  acid  decompose  one 
another,  what  substances  are  formed  ? 

141.  The  bromides  of  the  metals  closely  resemble  the 
chlorides ;   like  the  latter,  they  are  nearly  all  soluble  in 
water.     Silver,  lead,  and  mercurous  bromides  are  insoluble. 

To  a  solution  of  potassium  bromide  add  a  few  drops  of 
a  solution  of  silver  nitrate.  What  is  the  precipitate  ?  Write 
the  reaction. 

To  another  portion  of  the  potassium  bromide  add  a 
soluble  lead  salt.  To  a  third  portion  add  mercurous  nitrate 
solution. 

Explain  what  happens  in  each  case.  Compare  the  results 
with  similar  experiments  made  under  Hydrochloric  Acid. 

How  could  pure  hydrobromic  acid  gas  be  prepared  ? 
Read  descriptions  of  the  method  in  text-books.  Read  also 
about  the  properties  of  hydrobromic  acid,  and  notice  how 
it  resembles  and  how  it  differs  from  hydrochloric  acid. 
How  could  one  gas  be  distinguished  from  the  other? 
What  is  the  volumetric  composition  of  hydrobromic  acid  ? 

PREPARATION  OF  AMMONIUM  BROMIDE,  NH4Br. 

142.  A  flask  of  about  300  cubic  centimeters'  capacity  and  a 
50  cubic  centimeter  separating  funnel  are  required  (Fig.  27). 
Measure  1 10  cubic  centimeters  of  concentrated  ammonia 
solution  (sp.  gr.  .9)  into  the  flask,  and,  in  the  hood,  measure 
37  cubic  centimeters  of  bromine,  and  pour  into  the  funnel. 
Support  the  separating  funnel  in  a  clamp  stand ;  under 
it  place  the  flask  with  the  ammonia.     It  is  well  to  have 


66  LABORATORY  WORK  IN  CHEMISTRY. 

the  flask  standing  in  a  vessel  of  cold  water.  Allow  the 
bromine  to  run  into  the  ammonia,  drop  by  drop.  Shake 
the  flask  constantly  and  keep  it  cool.  Nitrogen 
gas  escapes,  fumes  of  ammonium  bromide  fill  the 
interior  of  the  flask.  Do  not  add  bromine  after 
the  alkaline  reaction  has  disappeared,  or  when  the 
liquid  becomes  yellow  add  more  ammonia.  The 
following  equation  represents  the  reaction  : 

4NH3  +  3Br=3NH4Br  +  N. 

Pour  the  contents  of  the  flask  into  a  large  por- 
celain evaporating  dish,  and  evaporate  to  dryness 
'  27'  on  the  water  bath.     The  salt  can  be  purified  by 
recry  stall  ization.     In  your  laboratory  book,  under  "  Nitro- 
gen," make  a  note  of  this  method  of  preparing  the  gas. 

PREPARATION  OF  POTASSIUM  BROMIDE,  KBr. 

143.  Dissolve  98  grams  of  ammonium  bromide  in  hot 
water,  and  add  100  grams  of  acid  potassium  carbonate. 
Heat  to  boiling,  and  continue  heating  until  the  odor  of 
ammonia  disappears.  Then  evaporate  and  crystallize  the 
potassium  bromide.  The  action  is  thus  represented  : 


Preserve  the  crystals  in  a  specimen  tube. 

IODINE,  I. 

144.  Preparation  of  Iodine.  —  Mix  about  one  gram  of 
potassium  iodide  with  about  twice  its  weight  of  manganese 
dioxide.  Put  the  mixture  in  an  evaporating  dish,  add  some 
dilute  sulphuric  acid,  stir  to  make  a  pasty  mass.  Place  a 
clean  funnel  over  the  dish  and  then  heat  it  gently.  Grad- 


HYDRIODIC  ACID. 


ually  the  funnel  becomes  filled  with  the  colored  vapor  of 
iodine,  and  grayish  black  crystals  are  deposited  on  the  cold 
sides  of  the  funnel.  Interpret  the  reaction. 

145.  Melt  two  or  three  crystals  of  iodine  in  a  small  test 
tube  and  pour  the  vapor  upon  a  sheet  of  white  paper.     Do 
you  obtain  crystals  or  an  amorphous  powder  ? 

146.  Solubility  of  Iodine  in  Solvents.  —  Make  solutions 
of  iodine  in  water,  in  alcohol,  and  in  a  water  solution  of 
potassium  iodide.     Use  small  quantities  in  test  tubes. 

147.  Action  of  Iodine  upon  Starch.  —  Make  some  starch 
paste  by  boiling  water  in  a  beaker  and  adding  to  it  a  small 
quantity  of  powdered  starch  mixed  with  cold  water ;   boil 
the  paste,  stirring  vigorously  with  a  glass  rod.     After  cool- 
ing, add  a  drop  of  the  water  solution  of  iodine  to  some  of 
the  starch  paste.     What  change  takes  place  ?     Add  starch 
paste  to  a  dilute  water  solution  of  potassium  iodide.     Is 
there  any  change  of  color  ?    Add  now  to  the  last  a  drop  or 
two  of  chlorine  water.     What  takes  place  ?    Does  chlorine 
alone  form  a  blue  compound  with  starch  ? 

Read  about  the  occurrence,  preparation,  properties,  and 
uses  of  iodine  in  the  text-books.  How  does  the  specific 
gravity  of  iodine  vapor  compare  with  that  of  hydrogen  ? 

HYDRIODIC  ACID,  HI. 

148.  Treat  a  few  crystals  of  potassium  iodide  with  concen- 
trated sulphuric  acid  in  a  test  tube.     What  do  you  notice  ? 
Hydriodic  acid  is  more  unstable  than  hydrobromic  acid.     In 
the  presence  of  the  sulphuric  acid  it  is  decomposed ;  free 
iodine  and   the  reduction  products  of   sulphuric  acid  are 
formed.     What  reduces  the  sulphuric  acid  ? 


68  LABORATORY   WORK   IN   CHEMISTRY. 

How  can  pure  hydriodic  acid  be  prepared  ?  What  are 
its  properties  ? 

149.  Iodides  of  the  Metals  Resemble  the  Chlorides  and 
Bromides.  —  Take  three  test  tubes  containing  dilute  solu- 
tions of  potassium  iodide ;  add  silver  nitrate  to  one,  lead 
nitrate  to  another,  mercurous  nitrate  to  the  third.     Explain 
what   takes   place   in   each   case,  and    compare  with   the 
similar  experiments  under  bromides  and  chlorides. 

How  can  iodides  be  distinguished  from  chlorides  and 
bromides  ?  ^ 

What  is*  the  volumetric  composition  of  hydriodic  acid  ? 
Compare^ith  hydrochloric  and  hydrobromic  acids  ? 

9  -"          * 

FLUORINE,  F,  AND  HYDROFLUORIC  ACID,  HF. 

Read  about  the  preparation  and  properties  of  fluorine. 
What  is  the  chief  source  of  fluorine  compounds? 

150.  Preparation  of  Hydrofluoric  Acid.  —  Cover  the  sur- 
face of  a  piece  of  glass  with  paraffine,  and  with  a  pointed 
instrument,  such  as  the  end  of  a  file,  scratch  letters  or 
figures  through  the  paraffine  so  as  to  leave  the  glass  ex- 
posed where  the  scratches  are  made.     In  a  lead  dish  put 
5  or  6  grams  of  powdered  fluorspar,  and  pour  on  it  enough 
concentrated  sulphuric  acid  to  make  a  thick  paste  when 
stirred  up  with  an  iron  wire.     Put  the  glass  with  the  paraf- 
fined side  downward  over  the  vessel  containing  the  fluor- 
spar, and  let  it  stand  for  several  hours.     Take  off  the  glass, 
scrape  off  the  paraffine.     The  figures  which  were  marked 
through  the  coating  will  be  found  etched  in  the  glass. 

How  does  the  preparation  of  hydrofluoric  acid  compare 
with  that  of  hydrochloric  acid  ?  Write  the  equation.  What 
are  the  distinguishing  properties  of  hydrofluoric  acid  ? 


SULPHUR.  69 

SULPHUR,  S. 

151.  Properties  of  Sulphur.  —  Heat  small  pieces  of  roll 
sulphur  in  an  ignition  tube.      Heat  very  gently  at    first. 
Notice  that  it  melts,  forming  a  thin,  straw-colored  liquid. 
As  the  temperature  gets  higher  the  liquid  becomes  darker 
and  darker ;  after  a  time  it  gets  so  thick  that  the  tube  can 
be  turned  upside  down  without  danger  of  its  running  out. 
At  higher  temperatures  it  becomes  more  liquid  again,  and 
at  440°  it  boils.    Having  noticed  these  changes,  pour  some 
of  the  hot  sulphur  into  cold  water  and  examine  the  plastic 
sulphur  thus  obtained. 

152.  Monoclinic  or  Prismatic   Sulphur.  —  In  a  covered 
Hessian  crucible  or  small  beaker  heat  a  quantity  of  roll 
sulphur.     After  it   has  all  melted  let  it  cool  slowly,  and 
when  a  thin  crust  has  formed  on  the  surface  make  a  hole 
through  this  and  pour  out  the  liquid  part  of  the  sulphur. 
The  inside  of  the  vessel  will  be  lined  with  honey-yellow 
needles.     Take  out  a  few  crystals  and  examine  them.    Are 
they  brittle  or  elastic  ?    Note  their  color,  and  whether  they 
are  opaque,  transparent,  or  translucent.     Lay  the  crucible 
aside,  and  in  the  course  of  a  few  days  again  examine  the 
crystals.     What  changes  have  taken  place  ? 

153.  Rhombic  or  Octahedral  Sulphur.  —  Dissolve  2  or  3 
grams  of  roll  sulphur  in  5  or  10  cubic  centimeters  of  carbon 
disulphide.     Put  the  solution  in  a  watch  crystal,  and  allow 
it  to  evaporate  by  standing  in  the  air.     What  is  the  appear- 
ance of   the  crystals  ?     Are   they  dark  yellow  or   bright 
yellow  ?     Are   they  brittle  or   elastic  ?      State  in  tabular 
form  the   properties  of  the  allotropic  forms    of   sulphur. 
Which   is  the    more    stable    above    and    below   96°    C.  ? 


/O  LABORATORY   WORK  IN   CHEMISTRY. 

154.  In  an  ignition  tube  heat  sulphur  to  boiling.     Intro- 
duce into  the  vapor  a  strip  of  thin  sheet  copper,  or  hold 
a  narrow  strip  so  that  the  end  just  dips  into  the  boiling 
sulphur.    What  evidence  have  you  that  action  takes  place  ? 
What  is  formed  ? 

Where  is  sulphur  found  in  nature  ?  How  is  it  purified 
and  what  are  its  chief  uses  ?  Name  some  of  the  principal 
sulphur  compounds  found  in  nature. 

HYDROGEN  SULPHIDE,  H2S. 

155.  Preparation  from  Ferrous  Sulphide,  FeS.  —  Into  a 

flask  put  a  small  handful  of  pieces  of 
iron  sulphide,  FeS.  Close  the  flask  with 
a  stopper  provided  with  a  funnel  tube 
and  a  delivery  tube  (Fig.  28).  Pour 
dilute  sulphuric  acid  upon  the  iron  sul- 
phide. Pass  the  gas  into  water  in  a 
test  tube.  Some  of  it  dissolves,  and  the 
water  acquires  the  odor  of  the  gas. 
Collect  a  gas  jar  full  of  the  gas  by  down- 
ward displacement  as  in  the  case  of  chlo- 
rine. (Specific  gravity  1.178.)  Set  fire 
FIG.  28.  to  tne  gas  in  the  jar  What  are  the 

products  of  combustion  ?     Write  the  equation  representing 
the  formation  of  hydrogen  sulphide. 

156.  Pass  hydrogen  sulphide  successively  through  solu- 
tions of  lead  nitrate,  zinc  sulphate,  and  of  arsenic  chloride 
(prepared  by  dissolving  a  little  white   arsenic,  As2O3,  in 
dilute  hydrochloric  acid).     What  do  you  observe  in  each 
case?     Sulphides  of  lead,  zinc,  and  arsenic  are  formed. 
Endeavor  to  write  the  equation  of  reaction  in  each  case. 


SULPHUR  DIOXIDE.  71 


Can  sulphur  and  hydrogen  be  made  to  combine  directly  ? 
How  is  hydrogen  sulphide  employed  in  analytical  chem- 
istry ?  What  would  be  the  effect  of  passing  hydrogen 
sulphide  over  heated  iron  ?  What  takes  place  when  steam 
is  passed  over  heated  iron  ?  How  can  small  quantities  of 
hydrogen  sulphide  be  detected  ? 

SULPHUR  DIOXIDE,  S02. 

157,  Preparation  from  Copper  and  Sulphuric  Acid.  —  Put 
8  or  10  pieces  of  sheet  copper,  one  or  two  inches  long  and 
half  an  inch  wide,  into  a  glass  flask. 

Pour   20    to    30    cubic    centimeters 

of  concentrated  sulphuric  acid  upon 

it.     Close  the  flask  with  a  stopper 

provided  with  a  delivery  tube  (Fig. 

29).    Heat  the  flask  gently  and  not 

too  rapidly  when  the  gas  begins  to 

come  off.    Saturate  water  in  a  bottle 

with  the  gas.    Collect  one  or  two  dry 

jars   full  of  the  gas.     This   can   be 

readily  done   by  the  downward  dis-  FlG>  29- 

placement  of  air,  as  the  gas  is  more  than  twice  as  heavy 

as  air. 

By  what  other  methods  can  the  gas  be  obtained  ?  Write 
the  equations.  What  volume  of  oxygen  is  contained  in 
one  volume  of  sulphur  dioxide  ? 

158.  Notice    the    physical    properties   of  the   gas,   the 
appearance,  odor,  etc.     Apply  a  burning  match  to  the  gas 
in  the  jar.     Does  it  burn?     Does  it  support  combustion? 
Reserve  one  bottle  of  sulphur  dioxide  for  the  experiments 
under  sulphuric  acid. 


72  LABORATORY  WORK  IN   CHEMISTRY. 

159,  Bleaching  Action  of  Sulphur  Dioxide.  —  Burn  sul- 
phur in  a  gas  jar  or  under  a  bell  jar. 
Place  over  the  burning  sulphur  some  red 
flowers  (Fig.  30).  Allow  them  to  remain 
in  the  atmosphere  of  sulphur  dioxide. 
They  will  be  bleached. 

160,    To  a  solution  of  rose  aniline  add 
some  of  the  solution  of  sulphur  dioxide 
FIG>  3a         '    in  water  obtained   in    Experiment   157. 
What  is  the  effect? 

161.  Neutralize  some  of  the  solution  of  sulphur  dioxide 
with  caustic  soda.     Then  evaporate  it  to  dryness.     What 
salt  is  formed  ?     How  is  sulphur  dioxide  prepared,  and  for 
what  purposes  is  it  used  in  the  arts?     Can  it  be  easily 
liquefied  ? 

162.  General   Behavior   of   Sulphites.  —  Sulphites,   like 
carbonates,  are  decomposed  by  hydrochloric  acid.     Pour 
hydrochloric  acid  upon  solid  sodium  sulphite  in  a  tube. 
Notice  the  effervescence  and  the  odor  of  the  gas. 

Nascent  hydrogen  reduces  them,  and  hydrogen  sulphide 
is  evolved.  To  a  solution  of  sodium  sulphite  add  zinc 
and  hydrochloric  acid.  The  hydrogen  sulphide  can  be 
recognized  by  its  smell  and  its  blackening  action  on  paper 
moistened  with  a  solution  of  a  lead  salt. 

Barium  chloride  produces  a  white  precipitate  of  barium 
sulphite,  BaSO3,  soluble  in  hydrochloric  acid.  This  solu- 
tion, on  the  addition  of  chlorine  water,  yields  a  white 
precipitate  of  barium  sulphate,  the  sulphite  being  oxidized 
to  sulphate.  See  reactions  for  sulphuric  acid. 

How  can  sulphur  dioxide  be  made  to  unite  with  more 
oxygen  ?  What  are  the  chief  properties  of  sulphur  trioxide  ? 


SULPHURIC  ACID.  73 

SULPHURIC  ACID,  H2S04. 

163.  Preparation  from  Sulphur,  Saltpeter,  and  Water.  — 
Pour  a  few  drops  of  water  into  a  glass  jar  and  moisten  the 
sides  of  the  vessel,  then  introduce,  by  means  of  the  defla- 
grating spoon,  a  burning  mixture  of  sulphur  and  potassium 
nitrate.    Cover  the  jar,  and  after  letting  it  stand  some  time 
rinse  the  vessel  with  distilled  water,  pouring  the  water  into 
a  test  tube.    It  is  a  dilute  solution  of  sulphuric  acid.    Test 
it  with  litmus  paper  and  with  barium  chloride.     A  white 
precipitate   with   the   latter   reagent    consists    of    barium 
sulphate,  BaSO4,  and  shows  that  sulphuric  acid  has  been 
formed. 

164.  Preparation  from  Sulphur  Dioxide.  — To  a  jar  con- 
taining sulphur  dioxide  add  four  or  five  drops  of  concen- 
trated nitric  acid.     Shake  the  jar,  add  a  little  water,  and 
transfer   the    solution    to   a    test   tube    and    add    barium 
chloride.     Interpret  the  reaction. 

166,  Preparation  from  Sulphur  and  Nitric  Acid.  —  Boil  a 
very  small  quantity  of  sulphur  flowers  with  concentrated 
nitric  acid  in  a  test  tube ;  then  dilute  with  water  and  add 
a  few  drops  of  barium  chloride  solution.  What  is  formed, 
and  what  does  this  experiment  show  ? 

How  is  sulphuric  acid  manufactured  on  the  large  scale  ? 
Write  out  a  brief  account  of  the  essential  features  of  the 
process.  For  what  purposes  is  sulphuric  acid  used  in  the 
arts  ?  What  is  the  specific  gravity  of  concentrated  sul- 
phuric acid  ? 

166.   Properties  of  Sulphuric  Acid.     It  is  a  Strong  Acid. 
Add  one  drop  of  concentrated  sulphuric  acid  to  a  large 
REISER'S  LAB.  CHEM. — 6 


74  LABORATORY  WORK   IN  CHEMISTRY. 

beaker  full  of  water  and  test  the  solution  with  blue  litmus 
paper. 

167.  Heat  Evolved  when  Mixed  with  Water.  —  Into  30 
cubic  centimeters  of  water  in  a  beaker  pour  gradually  120 
grams  of  the  concentrated  acid,  stirring  the  mixture  with 
a  narrow,  thin  test  tube  containing  a  few  drops  of  alcohol, 
or  water.     Does  the  liquid  in  the  tube  boil  ?     Determine 
the   temperature   of   the    diluted    sulphuric    acid   with    a 
thermometer. 

168.  Into  a  test  tube  containing  concentrated  sulphuric 
acid  insert    a  splinter  of  wood.     Why   is    it   blackened  ? 
Place  about  5  grams  of  sugar  in  an  evaporating  dish  and 
pour   10  cubic  centimeters  of  concentrated  acid  upon  it. 
Heat  gently  and  stir  with  a  glass  rod.    Dip  a  glass  rod 
into   dilute    sulphuric   acid   and  draw  letters  on  ordinary 
writing  paper,  then  dry  the  paper  by  holding  it  near  the 
lamp.     What  do  these  experiments  show  ? 

169.  Four  73  cubic  centimeters  of  concentrated  sulphuric 
acid  into  27  cubic  centimeters  of  water  in  a  beaker.     When 
cold,  transfer  the  mixture  to  a  measuring  cylinder  and  read 
the  volume.     Is  it  greater  or  less  than  100  cubic   centi- 
meters ? 

What  inference  can  be  drawn  from  this  ? 

170.  Action    upon    Metallic    Oxides.  —  Dissolve    some 
powdered  copper  oxide  in  dilute  sulphuric  acid  with  the  aid 
of  heat.     Filter  the  solution,  then  evaporate  it  and  obtain 
crystals  of  copper  sulphate.     In  general,  what  is  formed 
when  an  oxide  dissolves  in  sulphuric  acid  ?    What  is  formed 
when  a  nitrate  is  heated  with  concentrated  sulphuric  acid  ? 
What  action  does   sulphuric   acid   have   upon   chlorides  ? 
Upon  metals  ? 


PHOSPHORUS.  75 

171.  General  Properties  of  Sulphates.  -  -  To  a  solution 
of  any  sulphate  add  a  solution  of  barium  chloride.  What 
is  the  change  ?  Is  barium  sulphate  soluble  in  water  and 
acids  ? 

To  a  solution  of  any  sulphate  add  a  solution  of  lead 
nitrate  or  acetate.  What  is  precipitated  ? 

Powder  a  little  of  the  sulphate  in  a  mortar,  mix  with  it 
sodium  carbonate,  and  heat  the  mixture  in  the  reducing 
flame  of  the  blowpipe  on  charcoal.  The  sulphate  is 
reduced,  and  sodium  sulphide,  Na2S,  is  formed.  This, 
when  moistened  with  water  on  a  silver  coin,  gives  a  black 
stain.  Treated  with  dilute  acids,  the  sodium  sulphide  gives 
off  hydrogen  sulphide.  This  is  an  exceedingly  delicate 
test  for  sulphur  in  any  of  its  compounds. 


PHOSPHORUS,  P. 

172.  Examine  a  stick  of  phosphorus  under  water,  try 
to  cut  it,  bend  it ;  observe  the  color.     Take  a  small  piece 
out  of  water ;  notice  that  it  fumes  in  the  air.     Dry  a  small 
piece  between  filter  paper,  and  dissolve  it  in  a  few  cubic 
centimeters  of  carbon  disulphide  in  a  test  tube.     Pour  the 
solution  upon  a  large  piece  of  filter  paper.     Hold  the  paper 
with  the  pincette,  and  wave  it  in  the  air.     The  disulphide 
evaporates  and  leaves  a  layer  of  phosphorus  finely  divided 
on  the  paper ;  this  takes  fire  spontaneously.     Take  care  in 
making  this  experiment,  and  pour  any  of  the  solution  that 
may  remain  into  the  sink. 

173.  Phosphorus  melts  easily,  and  takes  fire  when  heated 
slightly  above  its  melting  point.     Float  a  watch  crystal  on 
warm  water  in  a  water  bath  or  large  beaker.     Place  a  small 
piece  of  phosphorus,  not  larger  than  a  grain  of  wheat,  upon 


76  LABORATORY  WORK   IN   CHEMISTRY. 

it.  Heat  the  water.  The  phosphorus  melts,  and  soon  after 
takes  fire,  burning  with  a  bright  flame,  forming  phosphorus 
pentoxide,  P2O6,  and  a  small  quantity  of  red  phosphorus 
remains  behind. 

174.  Bring  together  in  an  evaporating  dish  a  very  little 
phosphorus  and  iodine.     What  takes  place  ?     What  is  the 
cause  of  the  heat  and  the  light  ?     What  other  examples 
have  you  had  of  the  direct  combination  of   elements  by 
simple  contact  ?     Of  direct  combination   of   elements   at 
elevated  temperatures  ? 

175.  Examine  red  phosphorus ;  try  its  solubility  in  car- 
bon disulphide.     Set  fire  to  some  of  it  by  heating  on  an 
iron  plate  in  the  hood.     What  is  formed  when  it  burns  ? 
Bring  a  little  red  phosphorus  in  contact  with  a  crystal  of 
iodine.     How  does  the  affinity  of  the  red  phosphorus  com- 
pare with  that  of  the  yellow  variety  ? 

In  what  forms  is  phosphorus  found  in  nature  ?  Is  it  a 
widely  distributed  element  ?  Of  what  importance  is  cal- 
cium phosphate  to  living  things  ?  How  is  the  element 
obtained  in  free  condition  ?  How  is  red  phosphorus  made 
from  yellow  phosphorus  ? 

176.  Preparation  of  Phosphine,   PH3,  or   Phosphor etted 
Hydrogen.  —  Drop  one  or  two  small  pieces  of  phosphorus 
into  a  strong  solution  of  caustic  potash    contained   in    a 
small  beaker.     Heat  the  beaker  on  a  sand   bath   in    the 
hood.     After  a  time  bubbles  of  phosphine   are   evolved ; 
these  take  fire  spontaneously  in  the  air.     Sometimes  beau- 
tiful smoke  rings  are  formed.     Potassium  hypophosphite 
is  also  formed. 


ARSENIC.  77 

Pure  phosphine  is  not  spontaneously  inflammable,  but 
some  liquid  phosphoretted  hydrogen,  P2H4,  is  formed  at 
the  same  time,  and  this  sets  fire  to  the  gas.  Phosphine  is 
poisonous.  How  many  compounds  of  hydrogen  and  phos- 
phorus are  there  ?  How  are  phosphorus  pentoxide  and 
phosphoric  acid  prepared  ?  In  which  experiments  have 
you  obtained  the  former  ? 

177.  General  Properties  of  the  Phosphates.  —  Make   a 
solution  of  disodium  hydrogen  phosphate  in  water,  and  to 
a  portion  of  it  add  a  few  drops  of  silver  nitrate  solution. 
A  light  yellow  precipitate  of  silver  phosphate,  Ag3PO4,  is 
formed. 

To  another  portion  of  the  phosphate  solution  add 
ammonia,  ammonium  chloride,  and  magnesium  sulphate 
solution.  A  white  crystalline  precipitate  of  ammonium 
magnesium  phosphate  is  formed.  This  reaction  is  used 
frequently  in  testing  for  phosphoric  acid. 

To  a  third  portion  of  the  sodium  phosphate  add  nitric 
acid,  and  then  a  solution  of  ammonium  molybdate  in  nitric 
acid.  A  yellow  color,  and  finally  a  yellow  precipitate,  is 
obtained.  The  reaction  is  hastened  by  warming  the  mixture. 

How  do  monobasic,  dibasic,  and  tribasic  acids  differ? 
Give  an  example  of  each  kind.  What  acids  are  obtained 
from  phosphoric  acid  by  heating  it  to  high  temperatures  ? 
Write  the  equations. 

ARSENIC,  As. 

178.  Examine  some  metallic  arsenic.     Notice  its  luster, 
and  that  it  is  quite  brittle.     Heat  a  little  of  it  in  an  igni- 
tion tube.     Does  it  melt  ?     Keep  on  heating  until  none  is 
left  on  the  bottom  of  the  tube ;  what  do  you  notice  on  the 
cold  parts  of  the  tube  ?     What  is  this  process  called  ? 


78  LABORATORY  WORK  IN  CHEMISTRY. 

Heat  a  little  of  the  element  on  charcoal  before  the  blow- 
pipe ;  it  takes  fire.  What  is  the  color  of  the  flame  and 
of  the  vapors  given  off  ?  Notice  the  peculiar  garlic-like 
odor.  What  is  formed  when  arsenic  burns  ?  Warm  a  few 
small  crystals  with  nitric  acid.  Do  they  dissolve  ? 

179.  Mix  together  about  equal  small  quantities  of  arsenic 
trioxide  and  finely  powdered  charcoal.     Heat  the  mixture 
in  a  small  dry  tube  closed  at  one  end.     The  arsenic  which 
is  set  free  will  be  deposited  on  the  walls  of  the  tube  in  the 
form  of  a  mirror.     Write  the  equation  representing  the 
reaction. 

What  compounds  of  arsenic  occur  in  nature  ? 
What  is  the  composition  of  the  substance  commonly 
called  "arsenic  "  or  "white  arsenic  ?  " 

180.  Heat  a  small  quantity  of  the  trioxide  in  a  dry  test 
tube.     What  happens  ?     Dissolve  some  in  water  to  which 
hydrochloric  acid  has  been  added,  and  conduct  hydrogen 
sulphide  into  the  solution.     What  is  formed  ?     Write  the 
equation. 

181.  Preparation    of    Arseniuretted    Hydrogen,    AsH3, 
or  Arsine.  —  Fit  up  a  flask  for  generating  hydrogen  from 
zinc  and  dilute  sulphuric  acid.     Connect  a  calcium  chloride 
U-tube  with  the  delivery  tube  of  the  flask,  and  connect 
with  the  (J-tube  a  piece  of  hard  glass  tubing  drawn  out  to 
a  small  diameter  at  one  end,  and  turned  upward.     Put  zinc 
in  the  flask,  and  pour  dilute  sulphuric  acid  upon  it.     When 
the  air  is  out  of  the  vessel,  light  the  hydrogen,  and  now 
add  slowly  a  little  of  a  solution  of  arsenic  trioxide,  As2O3, 
in  dilute  hydrochloric  acid.     What  change  takes  place  in 
the  color  of  the  flame  ?    Notice  the  fumes  given  off.   What 
are  they  ? 


ANTIMONY.  79 

182.  Hold  a  cold  piece  of  porcelain  —  for  example,  an 
evaporating  dish  —  in  the  flame  of  burning  hydrogen  and 
arsine.     Notice  the  appearance  of  the  spots  formed.     Heat 
the  hard  glass  tube  through  which  the  gas  is  passing  at 
one  point.     Just  in  front  of  the  heated  point  a  thin  layer 
of  metallic  arsenic  will  be  deposited.      This  is  called  a 
"  mirror  "  of  arsenic.     What  causes  this  deposit  ?     Minute 
quantities  of  arsenic  can  be  detected  by  the  experiments 
described.     It  makes  no  difference  in  what  form  of  combi- 
nation the  arsenic  is  put  into  the  flask.     This  experiment 
is  known  as  Marsh's  test  for  arsenic. 

183.  Arsenic  acid  and  the  arsenates  resemble  phosphoric 
acid  and  the  phosphates  in  their  behavior  towards  reagents. 
See  General  Properties  of  the  Phosphates,  Experiment  177. 
But  all  arsenic  compounds  are  precipitated  by  hydrogen 
sulphide  in  acidulated  solution.     Try  this  with  a  solution 
of  sodium  arsenate  acidified  with  hydrochloric  acid. 


ANTIMONY,  Sb. 

In  what  form  does  antimony  occur  in  nature  ? 
How  is  the  element  obtained  in  free  condition  ? 

184.  Examine  a  piece  of  metallic  antimony.  Notice  its 
white  color  and  brilliant  luster.  It  does  not  tarnish  in  the 
air  like  arsenic.  Heat  a  small  piece  on  charcoal  before 
the  blowpipe.  Does  it  burn  ?  Are  fumes  formed  ?  Is 
there  an  odor  as  in  the  case  of  arsenic  ?  Treat  a  few 
small  pieces  with  concentrated  nitric  acid  in  a  test  tube 
and  warm  gently  ;  a  white  powder  is  formed.  (Sb2O3  and 
HSbO3.) 

Antimony  is  insoluble  in  hydrochloric  acid. 


Bo  LABORATORY   WORK   IN   CHEMISTRY. 

185.  Dissolve  antimony  trioxide  in  concentrated  hydro- 
chloric acid.     Pour  a  few  drops  into  water ;  notice  the  for- 
mation of  a  white  precipitate,  SbOCl,  "  powder  of  algaroth." 

Into  another  portion  of  the  solution  pass  hydrogen  sul- 
phide ;  what  is  formed  ? 

186.  Preparation  of  Stibine,  or  Antimoniuretted  Hydro- 
gen, SbH3.  —  The  same  apparatus  and  the  same  method  as 
that  described  under  Arsine  is  used  for  preparing  stibine. 
When  hydrogen  is  being  given  off,  add  a  solution  of  tartar 
emetic   or   any   other  antimony  compound.      Notice  the 
change    in   the   appearance   of    the    flame    of    hydrogen. 
Introduce  a  piece  of  cold  porcelain  and  notice  the  anti- 
mony deposit.     It  is  darker  and  more  smoky  than  the 
arsenic  mirror.     Heat  the  hard  glass  tube  and  notice  the 
appearance    of    the    antimony    mirror   in    the    tube.      Is 
the  antimony  more  or  less  volatile  than  the  arsenic  ?     For 
methods  of  distinguishing  the  arsenic  from  the  antimony 
mirror,  consult  one  of  the  text-books  on  analytical  chem- 
istry. 

BORON,  B. 

187.  Preparation  of  Boric  Acid,   H3B03.  —  Dissolve    50 

grams  of  powdered  borax  in  150  cubic  centimeters  of 
water.  Filter  the  solution  if  it  is  not  clear,  and  then 
add  concentrated  hydrochloric  acid  until  the  reaction  is 
decidedly  acid.  Allow  the  mixture  to  stand  over  night. 
Crystals  of  boric  acid  separate.  Collect  the  crystals  in 
a  funnel,  wash  them  with  a  little  water,  allow  them  to 
drain.  Dry  them  by  pressing  between  filter  paper.  Dis- 
solve some  of  the  crystals  in  alcohol  in  a  porcelain  dish, 
and  set  fire  to  the  alcohol.  Note  the  color  of  the  flame, 
because  this  reaction  is  characteristic  of  boric  acid.  Write 


SILICON.  8 1 

the  equation  representing  the  formation  of  boric  acid  from 
borax.  Dip  turmeric  paper  into  an  aqueous  solution  of 
boric  acid,  then  allow  it  to  dry.  What  is  its  color  ? 

What  compounds  of  boron  occur  in  nature  ? 

How  is  the  element  obtained  in  the  free  state  ? 

188.  Make  a  bead  on  platinum  wire  by  melting  crystals 
of  boric  acid  in  the  loop.     Heat  with  the  blowpipe  until 
bubbles  of  steam  are  no  longer  given  off.     What  is  the 
compound  that  is  thus  obtained  ? 

What  is  the  composition  of  borax  ?  How  is  it  used  in 
distinguishing  certain  metal  oxides  from  one  another  ? 

SILICON,  Si. 

189.  To  a  solution  of  sodium  silicate  in  water  add  strong 
hydrochloric  acid.     Gelatinous  silicic  acid  is  precipitated. 
Evaporate  the  mixture  of  silicic  acid  and  sodium  chloride 
to  dryness  in  an  evaporating  dish.     Moisten  the  dry  mass 
with  hydrochloric  acid  and  again  evaporate.     Silicon  diox- 
ide, SiO2,  and  sodium  chloride  are  left.     Treat  with  water 
to  dissolve  the  latter,  and  filter  off  and  dry  the  silicon 
dioxide.     Explain  the  changes  that  have  taken  place. 

Mention  some  of  the  principal  varieties  of  silicon  diox- 
ide that  occur  in  nature.  In  what  other  forms  of  combi- 
nation is  silicon  found  ? 

How  can  the  element  be  isolated  ? 

What  element  does  silicon  resemble  in  many  respects  ? 

190.  Preparation  of    Silicon  Hydride,   SiH4.  —  Mix   to- 
gether intimately  a  saltspoonful  of  very  fine  silica  and 
the  same  quantity  of   powdered   magnesium.     Heat  the 
mixture  in  a  dry  test  tube.     Sometimes  the  action  is  very 
violent.     Magnesium  silicide  is  formed.    When  cold,  break 


82  LABORATORY  WORK   IN  CHEMISTRY. 

the  tube  and  drop  fragments  of  the  mass  into  dilute 
hydrochloric  acid  in  a  beaker.  Bubbles  of  SiH4  and 
hydrogen  are  given  off,  which  take  fire  when  they  come 
in  contact  with  the  air. 

191.  Preparation  of  Silicon  Tetrafluoride,  SiF4,  and  Hy- 
drofluosilicic  Acid,  H2SiF6.  —  Mix  intimately  50  grams  of 
calcium  fluoride  and  100  grams  of  sand  in  a  rnortar,  and 
pour  the  mixture  into  a  dry  flask.  Then  add  175  cubic 
centimeters  of  concentrated  sulphuric  acid,  and  shake  the 
flask  so  that  the  acid  comes  in  contact  with  all  parts  of 

the  mixture,  and  with  the  bottom 
and  sides  of  the  flask.  Arrange 
the  apparatus  as  in  Figure  31. 
The  delivery  tube  and  rubber 
connections  must  be  dry  inside. 
Heat  the  flask  gently  on  a  sand 
bath.  Silicon  tetrafluoride  is 
FIG  r  evolved.  Pour  200  cubic  centi- 

meters of  water  into  the  beaker, 

and  allow  the  funnel  to  dip  but  a  very  little  below  the 
surface  of  the  water,  and  stir  the  liquid  from  time  to  time 
with  a  glass  rod.  When  gas  is  no  longer  evolved,  filter 
the  solution  in  the  beaker.  The  jelly-like  mass  is  silicic 
acid.  Heat  some  of  it  in  an  evaporating  dish  and  obtain 
pure  silica,  SiO2.  The  clear  filtrate  contains  hydrofluo- 
silicic  acid,  H2SiF6.  Test  small  portions  with  solutions  of 
potassium  chloride  and  barium  chloride.  The  former 
gives  a  gelatinous  precipitate  of  K2SiF6 ;  the  latter,  a 
crystalline  precipitate  of  BaSiF6. 

Write  equations  representing  the  formation  of  silicon 
tetrafluoride  and  hydrofluosilicic  acid.  How  could  you 
show  that  glass  contains  silicon  ? 


POTASSIUM.  83 

CLASSIFICATION  OF  THE  ELEMENTS  AND  THE  PERIODIC 

LAW. 

In  your  notebook  make  a  list  of  the  metallic  and  non- 
metallic  elements  you  have  worked  with.  In  general,  how 
do  metals  differ  from  non-metals  ?  Is  it  possible  to  satis- 
factorily classify  all  the  elements  under  these  two  divi- 
sions ?  What  is  meant  by  a  natural  family  or  group  of 
elements  ?  Give  examples  of  such  families.  Is  there  any 
relation  between  the  atomic  weights  of  the  elements  in 
these  natural  families  ? 

What  is  the  system  of  classification  known  as  the  nat- 
ural or  periodic  system  ?  What  is  the  periodic  law  ?  Read 
in  the  text-books  about  the  periodic  recurrence  of  proper- 
ties in  elements,  when  they  are  arranged  in  the  order  of 
increasing  atomic  weight.  Make  a  list  of  properties  that 
vary  periodically.  Of  what  use  has  the  periodic  law  been 
in  the  discovery  of  new  elements  ?  Mention  several  ele- 
ments whose  existence  was  foretold  by  the  law. 

POTASSIUM,  K. 

Read  in  the  text-books  about  the  occurrence  of  potassium 
compounds  in  nature.  Of  what  importance  are  potassium 
compounds  to  plants  ?  Where  does  the  potassium  which 
plants  use  come  from  ?  When  plants  are  burned,  what 
becomes  of  the  potassium  compounds  ? 

192.  Heat  5  grams  of  acid  potassium  tartrate  in  a  small 
iron  evaporating  dish,  in  the  hood,  until  the  residue  is  white. 
[Any  other  potassium  salt  of  an  organic  acid  will  answer.] 
Extract  the  residue  with  water,  filter  the  solution,  and 
examine  it  by  means  of  red  litmus  paper.  Is  it  alkaline  ? 


84  LABORATORY  WORK  IN   CHEMISTRY. 

Examine  a  solution  of  potassium  carbonate  in  the  same 
way.  Is  it  alkaline  ?  Evaporate  the  solution  of  the  resi- 
due obtained  from  the  tartrate  to  dryness,  and  treat  what 
remains  with  hydrochloric  acid  in  a  test  tube.  Or  you  can 
treat  some  of  the  residue  obtained  by  heating  the  tartrate 
in  a  test  tube  with  hydrochloric  acid.  Test  the  gas  in  the 
tube  with  a  drop  of  limewater  on  the  loop  of  wire.  What 
does  this  experiment  show  ? 

How  is  the  metal  potassium  obtained  and  what  are  its 
chief  properties  ? 

193.  Take  a  lump  of  potassium  from  under  the  oil ;  put 
it  on  filter  paper.     Cut  off  the  outside  crust ;  notice  the 
metallic  luster  and  the  rapidity  with  which  the  metal  tar- 
nishes in  the  air.     Throw  a  small  piece  not  larger  than  a 
grain  of  wheat  upon  water.     What  takes  place  ?     What  is 
the  color  of  the  flame  ?     Is  the  solution  after  the  action 
alkaline  ?     Why  ? 

What  are  the  more  important  compounds  of  potassium  ? 

194.  Preparation  of  Potassium  Hydroxide,  KOH.  —  Dis- 
solve about  50  grams  of  potassium  carbonate  in  400  to  600 
cubic  centimeters  of  water.      Heat  to  boiling   in  a  large 
iron  vessel,  and  gradually  add  milk  of  lime  made  by  add- 
ing water  to  25  or  30  grams  of   quicklime.      During  the 
operation  the  mass  should  be  stirred,  and  water  should  be 
added  from  time  to  time  to  prevent  the  solution  from  getting 
too  concentrated.     Let  the  liquid  cool,  and  when  the  pre- 
cipitate has  settled,  decant  the  clear  liquid  into  a  bottle. 
This  is  a  solution  of   caustic  potash.      The  reaction  de- 
pends upon  the  fact  that  calcium  carbonate  is  insoluble, 
and  potassium  carbonate  and  hydroxide  are  soluble.    Write 
the  equation. 


POTASSIUM.  85 

195.  Examine  some  solid  potassium  hydroxide.     Allow 
a  small  piece  to  lie  exposed  to  the  air  for  several   days. 
What  action  does  the  air  have  upon  it  ?     What  is  formed 
when  its  solution  is  treated  with  hydrochloric,  nitric,  and 
sulphuric  acids  ?     How  many  salts  can  it  form  with  sul- 
phuric acid  ?     What  are  their  formulas  ? 

Under  what  conditions  is  saltpeter  formed  ?  Describe 
the  saltpeter  plantations.  How  is  saltpeter  used  in  mak- 
ing nitric  acid  ?  In  making  sulphuric  acid  ? 

196.  Mix  together  20  grams  of  saltpeter,  40  grams  of 
charcoal  powder,  and  3  grams  of  sulphur  flowers.     Put  the 
mixture  on  an  iron  plate  in  the  hood,  and  push  a  lighted 
burner  under  the  plate.     Stand  at  some  distance.     Give 
a  brief  acccunt   of  the  manufacture   of  gunpowder,  and 
explain  its  action  as  an  explosive. 

Nearly  all  the  potassium  compounds  are  soluble  in  water. 
The  acid  potassium  tartrate,  HKC4H4O6,  and  the  double 
chloride  of  potassium  and  platinum,  and  potassium  per- 
chlorate  are  less  soluble  than  the  other  compounds.  Ad- 
vantage is  taken  of  this  fact  in  testing  for  the  presence 
of  potassium. 

197.  To  a  moderately  strong  solu- 
tion of  potassium  chloride  add  a  few 
drops  of  perchloric  acid.     The  white 
precipitate  is  potassium  perchlorate, 
KC1O4. 

To    50   cubic   centimeters    of    the 
potassium    chloride    solution   add   a  FIG.  32. 

solution  of  tartaric  acid,  or  of  sodium 

bitartrate ;    allow  it  to  stand  after   stirring  with  a  glass 
rod.     A  white  crystalline   precipitate   of   acid   potassium 


86  LABORATORY  WORK   IN   CHEMISTRY. 

tartrate  is  formed,  HKC4H4O6.  Wash  a  little  of  the  pre- 
cipitate with  alcohol,  and  test  its  solubility  in  hot  water. 
What  is  its  reaction  towards  litmus  paper?  Heat  some 
to  a  high  temperature  in  an  evaporating  dish.  What  is 
formed  ? 

To  a  drop  of  potassium  chloride  solution  add  several 
drops  of  platinum  chloride  solution.  A  yellow  precipitate, 
K2PtCl6,  is  formed,  soluble  in  hot  water,  but  insoluble  in 
alcohol  and  ether. 

Dip  a  platinum  wire  into  the  solution  of  a  potassium 
salt,  and  bring  it  into  the  Bunsen  flame.  See  Fig.  32. 
Notice  the  color  imparted  to  the  flame.  Look  at  the 
flame  through  blue  glass.  This  is  a  delicate  test  for 
potassium  compounds. 

PREPARATION   OF   POTASSIUM    IODIDE,   KI. 

^  198.  To  3  grams  of  fine,  clean  iron  filings  suspended 
in  25  cubic  centimeters  of  water  in  a  well-cooled  flask, 
add  slowly  12^  grams  of  iodine.  A  slight  excess  of  iron 
must  be  used.  Filter  after  standing  some  time,  wash 
the  filter,  and  to  the  filtrate  add  2\  grams  of  iodine.  Dis- 
solve 8.2  grams  of  dry  potassium  carbonate  in  boiling 
water,  and  to  this  boiling  solution,  contained  in  an  evapo- 
rating dish,  add  the  solution  of  iodine  and  iron.  Filter 
off  the  precipitate  that  forms.  If  the  filtrate  is  not 
colorless,  more  potassium  carbonate  must  be  added  to  it. 
Evaporate  the  colorless  solution  of  potassium  iodide  and 
obtain  the  crystallized  salt. 


Pure  potassium  iodide  is  colorless,  and  when  its  solution 
is  acidified  with  sulphuric  acid,  it  must  not  turn  starch 
paste  blue.  Preserve  the  crystals  in  a  specimen  tube. 


SODIUM. 


PREPARATION   OF    POTASSIUM    CHLORATE,    KC103. 

199.  Dissolve  50  grams  of  potassium  carbonate  in  the 
smallest  possible  quantity  of  hot  water.  Conduct  chlorine 
gas  into  the  boiling  solution 
until  it  is  no  longer  alkaline. 
The  chlorine  gas  must  be  washed 
by  passing  through  water  before 
it  is  conducted  into  the  potas- 
sium carbonate  solution,  other- 
wise manganese  chloride  will 
be  carried  over,  and  pink  per- 
manganate will  be  formed.  An 
apparatus  arranged  on  the  plan 
of  that  shown  in  Fig.  33  may  be 
used.  Dilute  the  solution  of 
the  chlorate  with  hot  water  to  a 
volume  of  100  cubic  centimeters  ;  filter,  and  allow  the  salt 
to  crystallize.  Use  turmeric  paper  in  testing  the  solution. 

3  K2C03  +  3  C12 = 5  KC1  +  KC1O3 + 3  CO3. 
Preserve  the  crystals. 


SODIUM,   Na. 

What  compounds  of  sodium  are  found  in  nature  ? 
is  the  free  element  obtained  ?     Write  the  equation. 


How 


200.  Examine  a  piece  of  metallic  sodium;  cut  off  the 
outside  crust ;  notice  the  metallic  luster.  Throw  a  small 
piece  upon  water.  What  takes  place  ?  Throw  a  small 
piece  upon  filter  paper  floating  on  the  water.  Why  does 
the  hydrogen  take  fire  in  this  case  ?  After  the  action,  test 
the  water  with  red  litmus  paper.  Explain  why  it  is  alka- 
line. Compare  with  potassium. 


88  LABORATORY   WORK   IN  CHEMISTRY. 

What  are  the  chief  compounds  of  sodium  ? 
How  is  sodium  hydroxide  obtained  ? 
Describe  briefly  the  processes  of  making  sodium  carbo- 
nate from  salt.     For  what  purposes  is  soda  used  ? 

201.  The  carbonates  of  sodium  and  potassium  are  soluble 
in  water,  the  carbonates  of  all  other  metals,  excepting  those 
of   rubidium  and  caesium  and  some  acid   carbonates,  are 
insoluble ;  hence  sodium  carbonate  precipitates  solutions 
of  almost  all  other  metals. 

To  a  solution  of  calcium  chloride  add  a  few  drops  of  a 
solution  of  sodium  carbonate.  What  is  the  white  pre- 
cipitate that  forms  ? 

Treat  a  solution  of  zinc  chloride  with  sodium  carbonate. 
What  is  formed  in  this  case  ? 

To  a  solution  of  ferric  chloride  add  sodium  carbonate. 

Write  the  equation  representing  the  reaction  in  each 
case. 

202.  Sodium  salts  are  almost  all  soluble  in  water.     The 
most  characteristic  reaction  is  the  bright  yellow  color  which 
all  sodium  compounds  impart  to  the  Bunsen  flame.     Dip  a 
clean  piece  of  platinum  wire  into  a  solution  of  a  sodium 
salt,  and  then  bring  it  into  the  non-luminous  gas  flame. 
Notice  the  intense  yellow  color.     It  is  not  seen,  however, 
when  viewed  through  a  blue  glass.     It  is  thus  possible  to 
distinguish  potassium  salts  when  mixed  with  sodium  salts. 

203.  Preparation  of  Pure  Sodium  Chloride.  —  Make   a 
saturated  solution  of  common  salt,  50  grams  in  150  cubic 
centimeters  of  water,  filter  the  solution,  and  pass  into  it 
hydrochloric  acid  gas.      Allow  the  gas  to   pass    into  the 
solution  through  a  small  funnel  which   just   touches  the 
surface  of  the  liquid.      See   Fig.    16.      Sodium   chloride 


AMMONIUM   SALTS.  89 


separates  out.  Magnesium  chloride  and  other  saline 
impurities  remain  in  solution.  Decant  the  acid  solution, 
wash  the  residue  by  decantation  with  50  cubic  centime- 
ters of  cold  water,  allow  it  to  drain,  then  heat  to  dry- 
ness  with  constant  stirring  in  a  porcelain  dish.  Preserve 
the  pure  salt  in  a  specimen  tube. 

THE  SPECTROSCOPE. 

•  204.  Examine  the  construction  of  the  spectroscope. 
Upon  what  principle  is  it  based  ? 

Observe  the  continuous  spectrum  of  the  luminous  flame 
of  the  Bunsen  burner. 

Observe  the  spectra  given  by  sodium  chloride,  potassium 
chloride,  and  lithium  chloride,  by  dipping  clean  platinum 
wires  into  solutions  of  these  salts,  and  then  bringing  them 
into  the  non-luminous  flame.  Compare  the  colored  bands 
obtained  with  the  colored  plates  in  one  of  the  text-books 
on  chemistry. 

Notice  also  the  spectra  given  by  barium,  strontium,  and 
calcium  chloride  solutions. 

How  can  the  spectroscope  be  used  in  examining  sub- 
stances of  unknown  composition  ? 

AMMONIUM  SALTS. 

Ammonia  unites  directly  with  acids  to  form  salts  that  in 
many  ways  resemble  the  salts  of  potassium  and  sodium. 

205.  Place  near  each  other  two  vessels,  one  containing 
strong  hydrochloric  acid,  the  other  strong  ammonia.  Blow 
across  the  top  of  these  vessels.  Explain  what  you  see. 

Neutralize  dilute  hydrochloric  acid  with  a  dilute  solution 
of  ammonia,  and  evaporate  the  solution  to  dryness  on  the 
REISER'S  LAB.  CHEM. — 7 


90  LABORATORY   WORK   IN   CHEMISTRY. 

water  bath.     Compare  the  white  salt  obtained  with  potas- 
sium and  sodium  chlorides. 

206.  Heat  in  a  dry  test  tube  one  or  two  crystals  of  am- 
monium chloride.     Does  it  sublime,  and  what  is  deposited 
in  the  cold  parts  of  the  tube  ?     The  ammonium  salts  of  all 
volatile  acids  are  volatile,   sometimes  they  can   be   sub- 
limed, as  in  the  case  of  the  ammonium  chloride,  otherwise 
they  are  decomposed,  as  in  the  case  of  the  ammonium  salts 
of  the  oxygen  acids.     The  salts  of  non-volatile  acids  lose 
their  ammonia  when  fused.      Try   this  with   ammonium 
phosphate.     What  is  formed  when  ammonium  nitrate  and 
ammonium  nitrite  are  heated  ? 

207.  To  a  solution  of  any  ammonium  salt  .add  caustic 
soda  and  then  heat  the  mixture.     What  is  given  off  ?     The 
minutest  trace  of  an  ammonium  compound  can  be  detected 
by  mixing  the  substance  with  dry  calcium  hydroxide  in  a 
small  beaker,  and  covering  the  beaker  with  a  watch  glass, 
on  the  under  side  of  which  a  strip  of  moist  red  litmus 
paper,  or  moist  turmeric  paper  has  been  stuck.     Try  this 
with  a  small  quantity  of  sal  ammoniac. 

208.  The  acid  ammonium  tartrate  and  the  double  chlo- 
ride of  ammonium  and  platinum,  like  the  corresponding 
potassium  compounds,  are  soluble  with  difficulty  in  water. 

To  a  solution  of  an  ammonium  salt  add  acid,  sodium 
tartrate.  What  is  the  formula  of  the  precipitate  ? 

To  a  concentrated  solution  of  ammonium  chloride  add  a 
few  drops  of  a  solution  of  platinum  chloride.  The  pre- 
cipitate is  insoluble  in  alcohol  and  ether.  What  is  the 
reaction  ?  How  could  the  ammonium  platinum  chloride 
be  distinguished  from  the  potassium  platinum  chloride  ? 


CALCIUM.  91 

VOLUMETRIC  COMPOSITION  OF  AMMONIA. 

209.  Fill  a  graduated  tube,  closed  at  one  end  and  pro- 
vided with  a  stopcock  of  large  bore  at  the  other  end,  with 
water.  Invert  the  tube  over  water  in  the  pneumatic 
trough,  and  displace  the  water  in  the  tube  by  passing  in 
a  rapid  current  of  chlorine  free  from  air.  Close  the  stop- 
cock, remove  the  tube  from  the  water,  and  pour  into  the 
open  end  10  cubic  centimeters  of  concentrated  ammonia. 
Open  the  stopcock  cautiously,  and  allow  most  of  the 
ammonia  to  gradually  flow  down  into  the  part  containing 
the  chlorine.  Now  run  in  dilute  sulphuric  acid,  in  the 
same  way  in  which  the  ammonia  was  run  in,  to  more  than 
neutralize  the  excess  of  ammonia.  Next  fill  the  tube 
above  the  stopcock  with  water,  close  the  opening  with 
the  thumb,  and  invert  over  water  in  a  tall  cylinder.  Open 
the  stopcock,  and  lower  the  tube  until  the  gas  is  under 
atmospheric  pressure.  Measure  the  volume  of  nitrogen. 
Three  volumes  of  chlorine  should  liberate  exactly  one 
volume  of  nitroen. 


Instead  of  the  tube  with  a  glass  stopcock  of  large  bore, 
an  ordinary  glass  tube  closed  at  one  end  and  divided  into 
three  parts,  by  means  of  rubber  bands  on  the  outside,  may 
be  used.  This  is  filled  with  chlorine,  and  while  it  is  still 
in  the  pneumatic  trough,  it  is  closed  with  a  tightly  fitting 
rubber  stopper  carrying  a  small  dropping  funnel.  Then 
proceed  as  before. 

CALCIUM,  Ca. 

What  are  the  chief  compounds  of  calcium  found  in 
nature  ?  Read  in  the  text-books  about  the  method  used  in 
obtaining  metallic  calcium  and  the  properties  of  the  metal. 


Q2  LABORATORY   WORK   IN   CHEMISTRY. 

210.  Preparation  of  Calcium  Chloride  from  Calcium  Car- 
bonate.—  Dissolve  10  to  20  grams  of  limestone  or  marble 
in  ordinary  hydrochloric  acid.     Evaporate  the  solution  to 
dryness.     Heat  over  the  free  flame  to  drive  off  the  water  of 
crystallization.     Expose  a  few  pieces  of  the  residue  to  the 
air.     Does  it  become  moist  ?     In  which  experiments  has 
calcium  chloride  been  used  and  for  what  purpose  ? 

211.  Preparation    of    Calcium    Oxide,  or    Quicklime.  — 

Heat  a  small  piece  of  marble  on  charcoal  to  a  high  tem- 
perature with  the  blowpipe  flame.  When  cold,  bring  it 
upon  turmeric  or  red  litmus  paper,  and  moisten  with  a  drop 
of  water.  Explain  what  has  taken  place.  What  happens 
when  marble  is  heated  to  a  high  temperature  in  a  closed 
vessel  ?  What  is  the  law  of  dissociation  ? 

212.  To  40  or  50  grams  of   good  quicklime   add   100 
cubic  centimeters  of  water.     Afterwards  dilute  to  two  or 
three  liters  and  put  the  whole  in  a  wide-stoppered  bottle. 
Let  it  stand  several  hours,  then  pour  off  the  clear  solution 
of  calcium  hydroxide.     What  takes  place  when  some  of 
the  clear  solution  is  exposed  to  the  air  ?    When  the  breath 
from  the  lungs  is  passed  through  it  ?      Conduct  carbon 
dioxide  into  the  clear  limewater.     What  is  formed  ?    Keep 
up  the  current  of  CO2  until  the  precipitate  again  dissolves. 
What  is  in  solution  ?    Boil  the  clear  solution.    What  is  the 
precipitate  that  now  reappears  ? 

How  are  stalactites  and  stalagmites  formed  ? 

"^  213.  Heat  some  powdered  gypsum  to  about  200°  in  an 
air  bath.  Examine  what  is  left  by  mixing  with  a  little 
water  so  as  to  form  a  paste,  and  let  it  stand  to  see  whether 
it  will  harden.  How  is  plaster  of  Paris  made  ?  Upon 
what  does  the  hardening  depend  ?  What  causes  the 


DETERMINATION  OF  CARBON  DIOXIDE.  93 

"permanent  hardness  "  of  natural  waters  ?  How  can  it  be 
removed  ? 

214.    General  Reactions  of  Calcium  Compounds.  —  To  a 

solution  of  calcium  chloride  or  any  calcium  compound  add 
sodium  carbonate.  What  is  the  white  precipitate  that  is 
formed  ? 

Sulphuric  acid  precipitates  from  strong  solutions  white 
calcium  sulphate,  which  dissolves  in  a  large  excess  of  water. 
Try  it. 

Ammonium  oxalate  produces,  even  in  dilute  solutions, 
a  white  precipitate  of  calcium  oxalate,  insoluble  in  acetic 
acid  and  in  ammonia,  but  soluble  in  hydrochloric  and 
nitric  acids.  Calcium  chloride  introduced  into  the  non- 
luminous  gas  flame  gives  to  it  an  orange-red  color.  Try 
these  reactions. 


DETERMINATION  OF  CARBON  DIOXIDE  IN  CALCIUM 
CARBONATE. 

215.    Fit  up  an  apparatus  as  shown  in  the  figure. 
small  Erlenmeyer  flask  of  about   100  cubic  centi- 
meters' capacity  is  closed  with  a  stopper  provided 
with  a   calcium   chloride   tube,  and   a   tube   that 
extends  almost  to  the  bottom    of   the   flask.     A 
small  specimen  tube  or  test  tube  is  selected,  of 
such  size  that  it  can  easily  be  lowered  into  the 
flask  to  take  the  position  shown  in  Fig.  34.    Weigh 
out  accurately  .3  to  .5  grams  of  marble  in  a  watch 
crystal.    Transfer  the  marble  without  loss  to  the  flask,  and 
pour  on  it  a  few  cubic  centimeters  of  water  so  as  to  close 
the  opening  of  the  long  tube.    Fill  the  test  tube  two  thirds 
full  of  concentrated  hydrochloric  acid,  and  carefully  lower 


94  LABORATORY   WORK   IN   CHEMISTRY. 

it  into  the  flask  by  means  of  pincers.  Insert  the  stopper 
with  the  calcium  chloride  tube,  and  weigh  accurately  the 
whole  apparatus.  Now  tip  the  flask  so  that  the  acid  flows 
on  the  marble.  When  the  action  has  moderated,  warm 
gently.  Finally,  when  no  more  gas  is  given  off,  displace 
the  carbon  dioxide  in  the  apparatus  by  drawing  air  through 
by  suction  applied  at  the  end  of  the  calcium  chloride  tube. 
When  the  apparatus  is  perfectly  cold,  weigh  it.  The  loss 
in  weight  of  the  apparatus  is  the  weight  of  carbon  dioxide 
in  the  weight  of  marble  taken.  Calculate  the  per  cent. 

BARIUM,   Ba,  AND  STRONTIUM,   Sr. 

What  compounds  of  barium  and  strontium  are  found  in 
nature  ?  Barium  and  strontium  compounds  closely  resem- 
ble the  compounds  of  what  element  ? 

216.  The  carbonates,  sulphates,  and  oxalates  are  insol- 
uble in  water.    Try  this  by  adding  to  solutions  of  barium 
and  strontium  chlorides,  sodium  carbonate,  sulphuric  acid, 
and  ammonium  oxalate.     In  each  case  white  precipitates 
are  formed.     Write  the  equations. 

217.  There  are  differences  in  the  behavior,  however,  of 
these  elements,  which  can  be  used  in  detecting  their  pres- 
ence when  mixed  with  one  another  and  with  other  sub- 
stances.    Barium  compounds  moistened  with  hydrochloric 
acid  and  introduced  into  the  non-luminous  flame  impart 
a  green  color  to  it.     Strontium  compounds  under  similar 
circumstances  give  a  crimson  color. 

Hydrofluosilicic  acid  precipitates  solutions  of  barium 
salts,  but  not  strontium  or  calcium  solutions. 

Strontium  sulphate  solution  precipitates  barium,  but  not 
strontium  or  calcium  solutions. 


MAGNESIUM.  95 

Calcium  sulphate  solution  precipitates  both  barium  and 
strontium  solutions,  the  latter  after  standing  some  time, 
but  not  calcium  solutions. 

Potassium  chromate  precipitates  yellow  barium  chromate 
insoluble  in  acetic  acid.  The  same  reagent  does  not  pre- 
cipitate calcium,  and  strontium  is  precipitated  only  in 
concentrated  solutions.  The  strontium  chromate,  how- 
ever, is  soluble  in  acetic  acid.  Try  these  reactions. 

MAGNESIUM,  Mg. 

In  what  forms  does  magnesium  occur  in  nature  ? 

Read  in  the  text-books  about  the  preparation  of  the  free 
element.  What  happens  when  the  metal  is  heated  in  the 
air  ?  Is  it  possible  to  obtain  metallic  magnesium  by  the 
reduction  of  the  oxide  with  charcoal  at  high  tempera- 
tures ? 

218.  Magnesium  has  such  a  strong  affinity  for  oxygen 
that  it  will  reduce  oxides  that  are  not  reduced  by  carbon. 
Heat  together  in  a  dry  test  tube  as  much  of  a  dry  mixture 
of  magnesium  powder  and  fine  sand  as  would  fill  a  salt 
spoon.    The  action  is  rather  violent.    Transfer  the  product 
to  a  small  beaker,  and  add  dilute  hydrochloric  acid.     Silicon 
hydride  is  evolved,  which  bursts  into  flame  when  it  meets 
the  air.     Explain  what  has  taken  place.    What  action  does 
magnesium  have  upon  steam  at  high  temperatures  ? 

219.  Characteristic     Behavior     of     Magnesium     Com- 
pounds.—  To    a    solution    of     magnesium    chloride    add 
ammonia   solution ;    a   white    precipitate    of    magnesium 
hydroxide,  Mg(OH)2  is  formed.     Repeat  the  experiment, 
first  mixing  the  magnesium  chloride  with  an  equal  volume 
of  ammonium  chloride  solution.     No  precipitate  is  formed, 


96  LABORATORY  WORK  IN   CHEMISTRY. 

because  magnesium  hydroxide   is  soluble   in  ammonium 
salts. 


soluble  double  salt 

An  acidulated  magnesium  solution  is  not  precipitated  by 
ammonia,  because  ammonium  and  magnesium  salts  com- 
bine to  form  double  salts,  which  are  not  decomposed  by 
ammonia.  Try  the  experiment. 

220.  In    the    presence    of    ammonium    salts   and  free 
ammonia,  disodium  phosphate,  HNa2PO4,  precipitates  mag- 
nesium completely  as  magnesium  ammonium   phosphate, 
NH4MgPO4  •  6  H2O.    The  precipitate  is  at  first  voluminous, 
but  soon  becomes  crystalline.     In  a  very  dilute  solution  it 
forms  slowly.     The  precipitation  is  hastened  by  rubbing 
the  glass  with  a  rod.     The  precipitate  is  soluble  in  acids, 
and  is  reprecipitated  from  acid  solution  by  ammonia.    This 
is  the  reaction  most  frequently  used  in  testing  for  mag- 
nesium.    Try  it.     Write  the  equations  in  your  notebook. 

221.  Heat  a  piece  of   magnesite  intensely  before  the 
blowpipe.     What   change   takes   place  ?     Moisten  with  a 
drop   of   cobalt   nitrate   solution   and  heat  again.     What 
change  do  you  notice  ?     (Characteristic  Reaction  of  Mag- 
nesium Compounds.) 

ZINC,  Zn. 

What  are  the  most  important  zinc  ores  ?  How  is  the 
metal  obtained  from  its  ores  ?  How  is  it  utilized  in  the 
laboratory  ? 

222.  Heat  a  piece  of  zinc  on  charcoal  with  the  blowpipe 
flame.     Notice  the  crust  of  zinc  oxide.     What  is  its  color 


while  hot,  and  when  cold  ?  Moisten  the  oxide  with  cobalt 
nitrate,  and  heat  again.  What  color  has  it  now  acquired  ? 
This  is  a  very  characteristic  reaction  of  zinc  compounds. 

223.  Alkaline  hydroxides   precipitate  solutions  of   zinc 
salts.     The  zinc  hydroxide  formed  is  soluble  in  excess  of 
the   reagent,   zincates   being  formed.     Sodium    carbonate 
precipitates  basic  zinc  carbonate,  and  ammonium  sulphide, 
(NH4)2S,  gives  a  white  precipitate  of  zinc  sulphide,  ZnS, 
soluble  in  mineral  acids,  but  not  in  acetic  acid.     Convince 
yourself  of  these  facts  by  experiment,  and  write  the  equa- 
tions representing  the  reactions. 

CADMIUM,  Cd. 

The  metal  cadmium  resembles  zinc  in  appearance.  Its 
compounds  accompany  the  zinc  compounds  in  nature. 
How  is  cadmium  obtained  ? 

224.  Heat  a  small  piece  of  cadmium  on  charcoal  before 
the  blowpipe.     A  brown  incrustation  of  cadmium  oxide  is 
formed.     The  same  coating  is  obtained  by  heating  a  cad- 
mium compound  with  sodium  carbonate  on  charcoal. 

225.  Pass  hydrogen  sulphide  into  a  solution  of  cadmium 
salt  acidified  with  hydrochloric  acid.     A  yellow  precipitate 
of  cadmium  sulphide  is  formed.     This  is  a  characteristic 
reaction  for  cadmium.     Compare  with  zinc. 

MERCURY,  Hg. 

How  is  mercury  obtained  ?  What  are  its  physical  prop- 
erties, such  as  melting  point,  boiling  point,  specific  gravity, 
solubility  in  water  and  acids?  How  many  series  of  salts 
does  it  form  ?  How  many  oxides  of  mercury  are  there  ? 


98  LABORATORY  WORK   IN  CHEMISTRY. 

226.  Mercury  is  volatilized  from  all  mercury  compounds 
when  they  are  heated  with  soda.     Bring  a  mixture  of  dry 
mercuric  chloride  and  fused  soda   into  a  dry  glass  tube 
closed  at  one  end,  and  heat.     The  volatilized  mercury  is 
deposited  in  the  cooler  part  of  the  tube  as  a  mirror,  or  in 
drops. 

227.  Treat  a  small  drop  of  mercury  with  nitric  acid  that 
has  been  diluted  with  half  its  volume  of  water.     As  long 
as   mercury   is   in   excess,  what   is   formed  ?     Write   the 
reaction.     Compare  with  the   action  of   nitric  acid  upon 
copper  in  the  preparation  of  nitric  oxide.     How  is  mer- 
curic nitrate  obtained  ? 

228.  Put  in  a  test  tube  5  cubic  centimeters  of  mercurous 
nitrate,  and  in  another  5  cubic  centimeters  of  mercuric 
nitrate  solution  ;   add  slowly  hydrochloric  acid   to   each. 
Notice  the  difference  in  the  result.     Explain   what   has 
happened. 

In  the  same  way  treat  mercurous  and  mercuric  solutions 
with  caustic  soda.  Notice  the  different  behavior.  What 
is  formed  in  each  case  ? 

Try  also  the  effect  of  potassium  iodide  upon  mercurous 
and  mercuric  solutions.  Filter  the  red  mercuric  iodide, 
and  dry  it.  Then  into  a  clean  and  dry  test  tube  put  a 
small  amount  of  the  mercuric  iodide.  Warm  the  middle 
part  of  the  tube  until  you  cannot  bear  your  hand  upon  it. 
Then  heat  rather  gently  the  lower  part  of  the  tube  where 
the  mercuric  iodide  is.  Explain  what  you  see.  Try  the 
effect  of  rubbing  some  of  the  yellow  iodide  with  a  rod. 
What  other  substance  that  you  have  worked  with  exists  in 
two  crystalline  forms,  one  of  which  is  more  stable  at  ordi- 
nary temperatures  than  the  other  ? 


COPPER. 


99 


229.  To  solutions  of   mercurous   nitrate  and  mercuric 
chloride  add  hydrogen  sulphide.     What  is  formed  in  each 
case  ? 

To  mercurous  nitrate  add  a  solution  of  stannous  chloride. 
What  is  the  precipitate  ?  Pour  off  the  liquid,  and  boil  the 
gray  precipitate  with  hydrochloric  acid  until  globules 
appear. 

Try  the  action  of  stannous  chloride  upon  mercuric 
chloride.  At  first  a  white  precipitate  forms,  finally  a  gray 
one.  Explain  the  reaction.  (Characteristic  reaction.) 

•+ 

COPPER,  Cu. 

In  what  forms  does  copper  occur  in  nature  ?  How  is  it 
obtained  from  its  ores  ?  Is  the  metal  easily  fused  ?  What 
is  formed  when  it  is  heated  in  the  air  ? 

What  gas  is  obtained  by  the  action  of  nitric  acid  upon 
it  ?  What  gas  is  obtained  by  the  action  of  hot  sulphuric 
acid  upon  it  ?  What  is  brass  ?  What  is  bronze  ? 

How  many  oxides  of  copper  are  there  ?  What  are  their 
formulas  ? 

230.  Preparation  of  Cuprous  Oxide.  —  To  a  cold  dilute 
solution  of  copper  sulphate  add  a  few  cubic  centimeters  of 
a  solution  of  grape  sugar,  then  enough  caustic  potash  to 
dissolve  the  precipitate  which  at  first  forms.     Warm  the 
mixture  ;  a  yellow  precipitate  of  cuprous  hydroxide,  CuOH, 
at  first  forms.     Boil,  and  this  becomes  red,  cuprous  oxide, 
Cu2O,  being  formed. 

231.  Preparation  of  Cupric  Oxide.  —  Place  a  small  quan- 
tity of  copper  nitrate  in  an  evaporating  dish  or  crucible, 
and  heat  until  fumes  are  no  longer  given  off.     What  has 
taken  place  ? 


IOO  LABORATORY  WORK   IN   CHEMISTRY. 

Put  into  a  test  tube  a  colcl  solution  of  copper  sulphate, 
and  add  to  it  a  solution  of  caustic  soda.  What  is  the  pre- 
cipitate ?  Repeat  the  last  experiment,  but  instead  of  hav- 
ing cold  solutions,  heat  both  solutions  to  boiling,  and  mix 
while  hot.  What  is  formed  in  this  case  ? 

232.  Add  ammonia,  drop  by  drop,  to  a  solution  of  cop- 
per sulphate,  shaking  after  each  addition.     Describe  what 
you  observe.     This  is  a  very  delicate  test  for  minute  quan- 
tities of  copper  salts. 

233.  Potassium  ferrocyanide  gives  a  brown  precipitate 
with  copper  salts.     Add  one  drop  of  copper  sulphate  solu- 
tion to  a  test  tube  full  of  water,  then  a  drop  of  acetic  acid 
and  a  small  quantity  of  potassium  ferrocyanide.     Describe 
and  explain  the  change  that  you  observe.    This  is  a  very 
delicate  test  for  copper. 

234.  Saturate  10  cubic  centimeters  of  dilute  copper  sul- 
phate solution  with  hydrogen  sulphide.    Write  the  reaction. 

235.  Metallic  iron  and  zinc  precipitate  copper  in  the 
metallic  state  from  its  solutions.     Try  this ;  use  a  bright 
wire  nail  and  a  bright  piece  of  zinc.     What  becomes  of 
the  zinc  and  iron  ? 

SILVER,  Ag. 

In  what  forms  is  silver  found  in  nature  ?  How  is  silver 
extracted  from  its  ores  ?  Describe  the  cupellation  process 
of  separating  silver  from  other  metals.  What  are  the 
characteristic  properties  of  the  metal,  —  color,  malleability, 
specific  gravity,  etc.  ?  What  are  its  most  important  alloys  ? 

236.  Preparation  of  Pure  Silver  Nitrate  from  a  Silver 
Coin Dissolve  a  lo-cent  piece  in  about  10  cubic  centi- 


SILVER.  1 01 

meters  of  pure  concentrated  nitric  acid  diluted  with  an 
equal  volume  of  water.  Divide  the  solution  into  two 
portions.  Evaporate  one  portion  to  dryness  in  an  evapo- 
rating dish.  Heat  the  dry  residue  cautiously  until  it 
fuses,  and  the  blue  copper  nitrate  is  changed  to  black 
copper  oxide.  Do  not  heat  too  high,  or  the  silver  nitrate 
also  will  be  decomposed.  When  cold,  treat  the  residue 
with  water  and  filter  the  solution  of  silver  nitrate.  Explain 
the  changes  that  have  taken  place. 

To  the  second  portion  add  hydrochloric  acid  as  long  as 
a  precipitate  is  formed.  Filter,  and  wash  with  water  until 
the  wash  water  no  longer  has  an  acid  reaction.  Bring  the 
silver  chloride  into  a  porcelain  dish,  stir  it  with  water,  and 
add  caustic  soda ;  heat  to  boiling  and  add  grape  sugar  or 
milk  sugar  in  small  pieces.  This  reduces  the  silver  chloride 
to  metallic  silver.  Five  minutes'  boiling  and  a  piece  of 
sugar  as  large  as  a  hazel  nut  should  suffice.  Pour  off  the 
liquid,  decant  a  dozen  times  with  water,  until  the  water 
shows  no  alkaline  reaction.  Dissolve  the  residue  in  a 
little  warm  dilute  nitric  acid ;  filter  if  the  solution  is  not 
clear,  evaporate  to  dryness  on  the  water  bath.  Dissolve 
the  crystalline  residue  in  distilled  water.  Keep  the  solu- 
tion of  silver  nitrate  for  future  use. 

237.  General  Reactions  of  Silver  Compounds Hydro- 
chloric acid  and  soluble  chlorides  give,  with  silver  solu- 
tions (use  the  silver  nitrate  solution  that  you  have  pre- 
pared), a  white  curdy  precipitate  of  silver  chloride,  insoluble 
in  water  and  nitric  acid,  but  readily  soluble  in  ammonia, 
sodium  hyposulphite,  and  potassium  cyanide.  Caustic  soda 
produces  a  light  brown  precipitate  of  silver  oxide,  Ag2O, 
insoluble  in  excess  of  the  hydroxide,  but  soluble  in  ammo- 
nia. Try  these  reactions. 


IO2  LABORATORY  WORK  IN  CHEMISTRY. 

Hydrogen  sulphide  or  ammonium  sulphide  produces  a 
black  precipitate  of  silver  sulphide,  Ag2S,  soluble  in 
potassium  cyanide. 

Why  do  silver  coins  and  spoons  become  black  ? 

Try  the  action  of  solutions  of  sodium  carbonate,  sodium 
phosphate,  and  potassium  chromate  upon  silver  nitrate. 
Explain  by  means  of  an  equation  what  happens  in  each 
case. 

/  238.  All  silver  compounds,  when  fused  with  sodium 
carbonate  on  charcoal,  yield  a  metallic  bead  of  silver. 
Try  this,  using  silver  chloride. 

What  action  does  light  have  upon  silver  chloride?  How 
are  the  silver  compounds  used  in  photography  ? 

ALUMINIUM,  Al. 

Mention  some  of  the  principal  compounds  of  aluminium 
that  occur  in  nature.  How  is  the  metal  obtained  ?  What 
are  its  most  important  properties  ?  How  does  its  specific 
gravity  compare  with  that  of  the  metals  in  common  use  ? 

v  239.  Action  of  Acids  upon  Aluminium — Place  a  small 
piece  of  aluminium  foil  in  each  of  three  test  tubes,  and 
try  the  action  of  hydrochloric,  nitric,  and  sulphuric  acids 
upon  it.  What  happens  ? 

What  are  the  forms  of  aluminium  oxide  found  in  nature  ? 

What  is  obtained  by  heating  the  hydroxide  or  sulphate 
of  aluminium  ? 

240.  Heat  some  aluminium  sulphate  or  ammonium  alum 
upon  charcoal  before  the  blowpipe.  Notice  that  the 
oxide  becomes  luminous  when  highly  heated,  but  it  does 
not  melt.  Moisten  it  with  cobalt  nitrate  and  heat  again. 
What  is  the  color  of  the  mass  ?  (Characteristic  reaction.) 


ALUMINIUM.  IO3 

* 

Which  oxides  have  you  treated  with  cobalt  nitrate  in 
earlier  experiments,  and  what  changes  were  produced  ? 

241.  Preparation  and  Properties  of  Aluminium  Hydrox- 
ide.—  Aluminium  hydroxide  is  a  very  weak  base,' and 
towards  strong  bases  it  acts  like  an  acid.  To  a  solution 
of  alum  add  caustic  soda,  drop  by  drop.  A  gelatinous 
precipitate  appears  at  first,  but  finally  redissolves  when 
enough  caustic  soda  has  been  added.  (Reactions  ?)  Repeat 
the  experiment,  using  ammonia  instead  of  caustic  soda. 
Does  the  precipitate  redissolve  ? 

«•*  242.  Sodium  carbonate  and  ammonium  sulphide  added 
to  a  solution  of  alum  precipitate  aluminium  hydroxide; 
carbon  dioxide  is  set  free  in  one  case,  and  hydrogen  sul- 
phide in  the  other.  What  are  the  equations  representing 
the  reactions  ? 

Sodium  phosphate  precipitates  white  voluminous  alumin- 
ium phosphate,  A1PO4,  when  added  to  a  solution  of  alum. 
Try  these  reactions. 

243.  To  a  small  quantity  of  dilute  cochineal  solution 
add  an  equal  bulk  of  alum  solution,  and  then  add  ammonia. 
The  flocculent  aluminium  hydroxide  carries  down  the 
coloring  matter.  The  colored  precipitate  is  called  carmine 
lake.  How  is  alum  used  in  purifying  water  ? 

*  244.  Preparation  of  Alum.  —  Saturate  10  cubic  centi- 
meters of  water  with  powdered  aluminium  sulphate  in  a 
test  tube.  In  another  test  tube  prepare  a  saturated  solu- 
tion of  potassium  sulphate.  Mix  5  cubic  centimeters 
of  each  of  the  solutions.  Shake  the  tube  or  stir  with  a 
rod.  Filter,  and  examine  the  precipitate  with  a  lens.  Is 
it  crystalline  ? 


104  LABORATORY   WORK   IN   CHEMISTRY. 

TIN,  Sn. 

How  does  tin  occur  in  nature  ?  How  is  it  prepared  com- 
mercially ? 

245.  Examine  a  piece  of  tin  foil ;   notice  that  it  is  a 
white,  soft,  ductile  metal ;  that  it  does  not  change  in  dry 
or  moist  air.     Heat  a  piece  on  charcoal  before  the  blow- 
pipe ;  it  melts  very  easily.     At  a  higher  temperature  it 
burns,  forming  tin  dioxide,  SnO2.     Treat  tin  with  concen- 
trated nitric  acid;   a  white  compound,  metastannic  acid, 
SnO3H2,  is  formed.     Treat  tin  with  concentrated  hydro- 
chloric acid,  and  heat.     Stannous  chloride  is  formed. 

Mention  some  of  the  important  alloys  of  tin.  How  is 
tin  plate  made  ?  Tin  forms  two  oxygen  compounds,  stan- 
nous  oxide,  SnO,  with  basic  properties,  and  stannic  oxide, 
SnO2,  with  mainly  acid  properties. 

246.  Behavior  of  Stannous  and  Stannic  Compounds.  —  In 
each  of  two  test  tubes  put  I  cubic  centimeter  of  stannous 
chloride  solution  ;  dilute  each  with  10  cubic  centimeters  of 
water.     To  one  of   the  solutions   add  enough  potassium 
permanganate  solution  to  give  it  a  faint  pink  color.     Satu- 
rate the  contents  of  both  tubes  with  hydrogen  sulphide. 
Notice  the  different  colors  of  the  precipitates.     Explain 
what    has    taken    place.     Treat    both    precipitates   with 
ammonium  sulphide.     Do  they  dissolve  ? 

Mercuric  chloride  gives  with  stannous  solutions  at  first 
a  white,  then  a  gray,  precipitate.  Explain.  See  under 
Mercury,  Experiment  229.  Try  its  action  upon  stannic 
compounds.  What  happens  in  this  case  ? 

1    247.   All  tin  solutions,  when  treated  with  metallic  zinc, 
give  a  precipitate  of  metallic  tin  in  shining  laminae,  or 


LEAD.  105 

as  a  spongy  mass.  All  tin  compounds  fused  with  soda 
and  potassium  cyanide  on  charcoal  give  white  globules  of 
metallic  tin,  and  a  slight  white  coating  of  the  dioxide. 
Try  these  reactions. 

LEAD,  Pb. 

What  is  the  chief  ore  of  lead,  and  how  is  the  metal 
obtained  from  it  ? 

248.  Cut  a  piece  of  sheet  lead  with  a  pocket  knife ; 
notice   the   color   of   the   bright   metal,  its  softness   and 
ductility.     Heat  a  bit  on  charcoal.     It  melts  easily ;  when 
highly  heated,  it  burns.     What  is  formed?     Try  its  solu- 
bility in  nitric  acid,  and  in  hydrochloric  and  cold  sulphuric 
acids. 

249.  Metallic  Lead  from  Lead  Compounds.  —  Fuse  any 
lead  compound  with  soda  on  charcoal.     A  mal- 
leable globule  of  lead  and  a  yellow  coating  of 

lead  oxide  will  be  obtained. 

Suspend  a  piece  of  sheet  zinc,  as  in  Fig.  35, 
in  a  solution  of  lead  nitrate  or  acetate,  and  let 
it  stand  overnight.  The  lead  will  be  depos- 
ited in  crystallized  form.  Write  the  reaction. 

How  many  oxides  of  lead  are  there,  and  how  are  they 
obtained  ? 

250.  Treat  a  little  red  lead  with  dilute  nitric  acid.     Fil- 
ter the  solution.     What  is  the  brown  powder,  and  what  is 
in  solution  ?     Treat  some  of  the  brown  powder  with  con- 
centrated hydrochloric  acid.     What  is  formed  ? 

251.  Try  the  effect  of   the  following  reagents  upon  a 
solution  of  a  lead  salt,  such  as  lead  acetate,  and  explain 

REISER'S  LAB.  CHEM. — 8 


IO6  LABORATORY  WORK  IN  CHEMISTRY. 

what  happens  in  each  case  :  (i)  hydrochloric  acid  or  solu- 
ble chlorides,  (2)  sulphuric  acid  or  sulphates,  (3)  hydrogen 
sulphide,  (4)  potassium  chromate. 

THE  LAW  OF  SPECIFIC  HEATS. 

Do  different  substances  have  different  capacities  for 
heat  ?  What  is  meant  by  the  specific  heat  of  a  substance  ? 
How  can  the  specific  heat  of  solid  substances  be  deter- 
mined ?  Make  a  table  of  the  specific  heats  of  seven  or 
eight  metals,  then  multiply  the  specific  heat  of  each  metal 
by  its  atomic  weight.  Is  the  product  the  same  in  each 
case  ?  Suppose  that  the  temperature  of  63  Ib.  of  copper, 
1 08  Ib.  of  silver,  and  200  Ib.  of  mercury  were  raised  10 
degrees,  how  would  the  quantity  of  heat  required  to  raise 
the  temperature  of  each  metal  compare  with  that  required 
to  raise  the  temperature  of  the  others  ?  What  is  the  law 
of  specific  heats  ?  How  can  it  be  used  to  determine  atomic 
weights  ?  Do  the  specific  heats  of  the  elements  vary  peri- 
odically with  the  atomic  weights  ? 

BISMUTH,  Bi. 

This  element  resembles  lead  in  many  respects,  but  it  is 
also  related  to  arsenic  and  antimony. 

•j  252.  Examine  metallic  bismuth ;  notice  its  reddish  lus- 
ter. It  is  very  brittle.  Heat  a  fragment  on  charcoal 
before  the  blowpipe.  What  is  the  color  of  the  oxide  that 
is  formed  as  a  coating  on  the  charcoal  ? 

v  253.  Dissolve  a  small  quantity  of  bismuth  nitrate  in 
water  to  which  a  few  drops  of  nitric  acid  have  been  added. 
Into  a  portion  of  the  clear  solution  pass  hydrogen  sul- 


CHROMIUM.  lO/ 

phide  ;  browni^r black  bismuth  trisulphide  is  formed.  To 
another  portion  add  potassium  chromate ;  a  yellow  precipi- 
tate, Bi2(CrO4)3,  is  formed. 

Pour  a  few  drops  of  the  bismuth  solution  into  a  large 
quantity  of  water ;  a  white  precipitate  of  basic  bismuth 

salt  is  formed. 

* 

CHROMIUM,  Cr. 

What  mineral  is  the  chief  source  of  chromium  com- 
pounds ? 

j^Iow  is  potassium  chromate  obtained  from  it  ? 
Iftiat  are  the  two  important  oxides  of  chromium  ? 

Which  of  these  is  basic  and  which  acid  in  properties  ? 

254.  Behavior  of  the  Chromic  Salts.  —  In  these  the  chro- 
mium plays  the  part  of  a  base.     Use  a  solution  of  chrome 
alum  for  the  following  experiments.     Add  caustic  soda  to 
the  chromium  solution,  drop  by  drop.     What  is  formed  ? 
Dissolve  the  precipitate  in  an  excess  of  the  reagent,  then 
dilute  with  water  and  boil.     Explain  what  happens. 

Treat  portions  of  the  chromium  solution  with  ammonia 
and  ammonium  sulphide.  Interpret  the  reactions.  How 
do  these  reactions  of  chromium  compare  with  those  of 
aluminium  salts  ? 

255.  Conversion  of  Chromic  Salts  into  Chromates.  —  Fuse 
in  a  crucible  a  mixture  of  sodium  carbonate  and  saltpeter, 
and  to  the  fused  mass  add  a  small  quantity  of  chromic  oxide 
or  chrome  alum  ;  yellow  sodium  chromate  is  formed.     If  a 
solution  of  chromium  salt  is  made  alkaline  with  caustic  soda, 
and  if  an  oxidizing  agent,  such  as  chlorine  water,  potas- 
sium permanganate,  or  lead  peroxide,  is  added,  the  green 
solution  turns  yellow,  and  sodium  chromate  is  formed. 


IO8  LABORATORY   WORK   IN   CHEMISTRY.. 

256.  Behavior  of  Chromates.  —  Add  acra  to  the  yellow 
solution  of  sodium  or  potassium  chromate  until  it  turns 
red.     What  is  formed  ?     To  the  red  solution  add  an  alkali, 
NaOH  or  Na2CO3,  until  it  is* again  yellow.     Explain  the 
changes  that  have  taken  place.     To  a  solution  of  a  chro- 
mate add  barium  chloride.     Whatsis  the  precipitate  ? 

Treat  a  chromate  solution  with  a  soluble  lead  salt. 
What  is  formed? 

257.  Preparation  of  Chromium  Trioxide,  Cr03.  — Pour  20 
cubic  centimeters    of   a   saturated  solution  of   potassium 
dichromate  into  30  cubic  centimeters  of  concentrated  sul- 
phuric acid.    On  cooling,  red  needles  of  chromic  anhydride, 
CrO3,  separate.     Allow  the  crystals  to  subside,  pour  off 
the  liquid,  then  bring  some  of  the  crystals  upon  a  porous 
plate.     Try  the  solubility  of  some  of  the  crystals  in  water. 
Put  a  few  upon  filter  paper,  and  observe  what  effect  they 
have  upon  it.     Pour  a  drop  or  two  of  alcohol  upon  some 
of  the  crystals.      The  action  is  violent ;   sometimes  the 
alcohol  takes  fire.     What  is  the  green  substance  formed 
in  these  experiments  ? 

258.  Conversion  of  Chromates    into    Chromic  Salts. — 

Chromates  readily  give  up  oxygen,  and  are  reduced  to 
chromic  salts  when  treated  with  reducing  agents  in  the 
presence  of  an  acid.  To  a  solution  of  potassium  dichro- 
mate acidified  with  sulphuric  acid,  add  any  one  of  the  fol- 
lowing substances,  and  heat  until  the  color  becomes  green  : 
SO2,  H2S,  SnCl2,  alcohol,  oxalic  acid. 

What  is  formed  when  potassium  dichromate  is  heated 
with  concentrated  sulphuric  acid  ?  With  hydrochloric 
acid  ?  What  gas  escapes  iriPeach  case  ?  What  substances 
are  obtained  by  heating  ammonium  dichromate  ? 


MAN 


GANESE.  109 

1 


MANGANESE,  Mn. 


Which  compounds  of  manganese'  occur  in  nature  ? 
of  the  oxides  is  commonly  used  in  the  prepara- 
tion of  twp  important  non-metallic  elements  ? 

The  most  important  manganese  ..compounds  besides  the 
dioxide  are  the  manganou^  salts,  the  manganates,  and  per- 


manganates. 


259.  Properties  of  Manganous  Salts.  —  Make  a  solution 
of  manganous  sulphate,  and  add  to  a  portion  of  it  ammo- 
nium sulphide.  What  is  the  color  and  composition  of  the 

•*•  l 

precipitate  ?     Try  its  solubility  in  acids. 

To  another  portion  of  the  manganous  solution  add 
caustic  soda.  The  precipitate  rapidly  darkens  in  the  air. 
Explain  the  changes. 

To  manganous  sulphate  add  ammonia  ;  a  portion  of 
the  manganese  is  precipitated  as  Mn(OH)2.  Manganous 
hydroxide  acts  upon  ammonium  salts,  forming  soluble 
double  salts.  Thus  : 


HO  +  2  NH 


Compare  this  action  with  that  of  magnesium  hydroxide. 

These  double  salts  are  not  decomposed  by  ammonia, 
and  hence  when  ammonia  is  added  to  a  manganous  solu- 
tion containing  free  acids  or  ammonium  salts  no  precipi- 
tate is  formed.  But  on  standing,  oxygon  is  absorbed  from 
the  air,  and  brown  manganic  hydroxide  is  precipitated. 

2(2NH4Cl.MnCla)+4NH8+5H2O  +  O 
=  Mn2(OH)6  +  8NH4Cl. 


1 10  LABORATORY  WORK   IN   CHEMISTRY. 


m 

I    CHEMISTRY. 


260.  Preparation  of  Potassium  Manganate  and  Perman- 
ganate. —  Any  compound  of  manganese,  when  fused  with 
alkalis  in  the  presence  of  oxidizing  agents,  is  converted 
into  an  alkaline  manganate.     Heat  5  grams  of  KOH  with 
3  grams  of  KC1O3  in  a  crucible  slowly  until  the  mass  fuses. 
Add  5  grams  of  powdered  manganese  dioxide,  and  heat 
just  enough  to  keep  in  fusion.     Dark  green  potassium 
manganate  is  formed.     It  dissolves  in    a  small  quantity 
of  water,  with  a  green  color.     Pour  off  the  green  solution. 
Add  more  water  to  half  the  solution.    .^Brown  manganese 
perhydroxide,  MnO3H2  or  Mn(OH)4,  is  precipitated,  and 
the  color  of  the  solution  changes  to  a  beautiful  purple,  due 
to  the  formation  of  potassium  permanganate.     Add  acid  to 
the  rest  of  the  solution.     Explain  the  change. 

Owing  to  the  intense  color  of  potassium  manganate,  th£ 
least  trace  of  manganese  can  be  detected  by  fusing  the 
powdered  substance  with  soda  and  saltpeter.  The  small- 
est quantity  of  manganese  imparts  a  bluish  green  color  to 
the  fused  mass. 

261.  Oxidizing  Properties   of    Permanganates.  — ,To    a 

weak  solution  of  ferrous  sulphate,  acidified  with  sulphuric 
acid,  add,  drop  by  drop,  a  solution  of  potassium  perman- 
ganate, until  the  pink  color,  which  at  first  disappears,  be- 
comes permanent.  The  ferrous  sulphate  has  been  changed 
to  ferric  sulphate,  and  the  permanganate  to  potassium  and 
manganous  sulphates.  Try  to  write  the  equation.  Make 
a  dilute  solution  of  oxalic  acid.  Acidify  with  sulphuric 
acid,  heat  nearly  to  boiling,  and  add  permanganate  as 
before.  Write  the  equation. 

262.  Permanganic  Acid.  —  The  minutest  quantity  of  any 
manganese  compound,  when  heated  with  nitric  acid  and 
lead  peroxide,  can  be  detected  by  the  purple  color  of  the 


IRON.  MI 

permanganic  acid  which  is  formed.  Dilute  about  5  cubic 
centimeters  of  concentrated  nitric  acid  with  an  equal  volume 
of  water,  add  about  a  saltspoonful  of  red  lead  or  lead  per- 
oxide, heat  to  60°  or  70°,  and  now  add,  drop  by  drop,  a 
weak  solution  of  any  manganese  salt.  The  purple  color  is 
due  to  the  presence  of  permanganic  acid.  This  is  the 
most  delicate  manganese  test. 

263.  A  small  quantity  of  any  manganese  compound  in- 
troduced into  a  borax  bead,  and  heated  in  the  oxidizing 
flame,  gives  an  amethyst-colored  bead.  Heated  in  the  re- 
ducing flame,  the  bead  becomes  colorless.  Try  this. 


IRON,  Fe. 

Name  the  most  important  ores  of  iron.  For  what  pur- 
pose is  iron  pyrites  utilized  ?  Describe  briefly  the  method 
of  extracting  iron  from  its  ores.  What  are  the  differences 
between  cast  iron,  wrought  iron,  and  steel  ?  How  are 
wrought  iron  and  steel  obtained  from  cast  iron  ? 

264.  Iron  is  attracted  by  the  magnet ;   it   is   infusible 
before  the  blowpipe.     Finely  divided  iron,  such  as  pow- 
dered iron,  takes  fire  when  heated  in  the  air,  or  when  scat- 
tered into  a  flame.     An  oxide,  FegO^  is   formed.     With 
acids  it  evolves  hydrogen,  which,  owing  to  contamination 
with  hydrocarbons,  has  an  unpleasant  odor. 

How  many  oxides  of  iron  are  there  ?  Ho.w  many  classes 
of  iron  compounds  are  there  ? 

265.  Behavior  of  Ferrous  Compounds.  — Treat  iron  nails 
or  iron  filings  with  hydrochloric  acid   and  warm   gently. 
Ferrous  chloride,  FeCl2,  is  formed.     Keep  iron  present  in 


112  LABORATORY    WORK   IN  CHEMISTRY. 

excess  of  the  acid,  and  use  the  solution  for  the  following 
experiments.  To  a  few  cubic  centimeters  of  the  solution 
add  caustic  soda  solution  ;  a  white  precipitate,  FeO2H2, 
is  formed,  which  almost  instantly  changes,  acquiring  a 
dirty  green  and  ultimately  a  reddish  brown  color,  owing 
to  absorption  of  oxygen,  and  conversion  into  ferric  hydrox- 
ide, Fe(OH)3. 

Add  ammonium  sulphide  to  ferrous  chloride  solution  ;  a 
black  precipitate,  FeS,  ferrous  sulphide,  is  formed.  What 
is  the  reaction  ?  With  potassium  ferrocyanide,  K4Fe(CN)6, 
a  white  precipitate,  K2Fe(CN)6,  is  formed,  which  rapidly 
becomes  blue  by  oxidation  to  Fe4(Fe(CN)6)3,  Prussian 
blue. 

Add  potassium  ferricyanide  solution  to  ferrous  chloride ; 
Turnbull's  blue,  Fe3(Fe(CN)6)2,  is  formed.  Potassium 
sulphocyanide  with  ferrous  chloride  gives  no  precipitate  if 
ferric  salts  are  absent.  Try  it. 

266.  Conversion  of  Ferrous  into  Ferric  Compounds.  —  Boil 
10  cubic  centimeters  of  ferrous  chloride,  to  which  a  few 
drops  of  concentrated  nitric  acid  have  been  added,  until  the 
color  is  reddish  yellow.     Interpret    the   reaction.     Other 
oxidizing  agents,  such  as  potassium  permanganate,  bromine 
water,  potassium  chlorate,  and  hydrochloric  acid,  may  be 
used  instead  of  the  nitric  acid.     Reducing  agents,  such  as 
stannous  chloride,  zinc  and  hydrochloric  acid,  and  hydrogen 
sulphide,  convert  ferric  into  ferrous  compounds. 

267.  Behavior  of  Ferric  Compounds.  —  To  ferric  chloride 
solution  add  caustic  soda  or  ammonia.     What  is  formed  ? 
Compare  with  ferrous  compounds. 

Ammonium  sulphide  produces  black  FeS,  mixed  with 
sulphur. 


NICKEL  AND   COBALT.  113 

Potassium  ferrocyanide  gives  Prussian  blue, 
Fe,(Fe(CN)6)3. 

Potassium  ferricyanide  with  ferric  salts  in  the  absence 
of  ferrous  salts  changes  the  color  of  the  solution  to  reddish 
brown,  but  does  not  produce  a  precipitate. 

Potassium  sulphocyanide,  even  in  very  dilute  solutions, 
produces  a  blood  red  color.  (Delicate  reaction.) 

Try  all  these  tests  and  write  the  equations  representing 
the  changes. 

NICKEL,  Ni,  AND  COBALT,  Co. 

Nickel  and  cobalt  are  white  metals.  They  resemble 
iron.  Like  iron,  they  are  attracted  by  a  magnet.  They 
dissolve  slowly  in  hydrochloric  and  sulphuric  acids,  and 
readily  in  nitric  acid. 

268.  Make  solutions  of  nickel  and  cobalt  nitrates  in 
water.  Add  ammonium  sulphide  to  each  solution,  and 
explain  what  takes  place.  In  the  same  way,  try  the  effects 
of  ammonia,  caustic  soda,  and  potassium  cyanide  solutions 
upon  the  nickel  and  cobalt  solutions.  Notice  the  different 
behavior.  Add  potassium  nitrite  to  solutions  of  nickel 
and  cobalt  acidified  with  acetic  acid.  In  the  case  of  cobalt 
a  yellow  precipitate  is  obtained. 


114 


LABORATORY   WORK  IN   CHEMISTRY. 


TABLE  OF  THE  ELEMENTS 

WITH   SYMBOLS  AND   ATOMIC  WEIGHTS. 


NAME. 

SYMBOL. 

ATOMIC 
WEIGHT. 

NAME. 

SYMBOL. 

ATOMIC 
WEIGHT. 

Aluminium      .  . 

Al 

27. 

Nd 

I4.O.  c. 

Sb 

1  2O 

Nickel 

Ni 

C.8  7 

As 

7C.. 

Nitrogen  

N 

j°'  / 
14.  O3 

Barium     

Ba 

137- 

Osmium  

Os 

i^.vy^ 

IQO.8 

Bismuth   

Bi 

2O8.  Q 

o 

iyy.0 
1  6. 

B 

II. 

Palladium  

Pd 

106  6 

Bromine               .      . 

Br 

70.  qc 

Phosphorus  

p 

-31. 

Cadmium  

Cd 

112. 

Platinum  

Pt 

IQC,. 

Caesium  

Cs 

132.9 

Potassium  

K 

3Q.  11 

Calcium  

Ca 

4O. 

Praseodymium.  .  . 

Pr 

147.  c 

Carbon  

c 

12. 

Rh 

103. 

Cerium 

Ce 

I4.O.2 

Rubidium  

Rb 

85.5 

Chlorine  

Cl 

-2C.4.C 

Ruthenium  

Ru 

101.6 

Chromium  

Cr 

C.2.  1 

Sm 

I  C.O. 

Cobalt 

Co 

CO. 

Scandium    ... 

Sc 

44- 

Columbium 

Cb 

04. 

Selenium  

Se 

70. 

Copper.  .  , 

Cu 

63.6 

Silicon  

Si 

28.4 

Erbium  

Er 

166.3 

Silver  

Ag 

107.92 

Fluorine 

F 

10. 

Sodium 

Na 

23.OC, 

Gadolinium       .    . 

Gd 

ic.6.  i 

Strontium     

Sr 

87.65 

Gallium  

Ga 

*:>"•  * 

6q. 

Sulphur  

s 

32.06 

Germanium  

Ge 

72.3 

Tantalum  

Ta 

182.6 

Glucinum 

Gl 

Tellurium  

Te 

I2C,. 

Gold  

Au 

107.  3 

Tb 

J 
1  60. 

Hydrogen  

H 

I.  OO7 

Thallium  

Tl 

204.18 

Indium 

In 

1  1  3  7 

Thorium          . 

Th 

2^2.6 

Iodine                      . 

I 

126.85 

Thulium  

Tu 

I7O.7 

Indium  

Ir 

IQ3.  1 

Tin         

Sn 

119. 

Iron  

Fe 

S^. 

Ti 

48. 

Lanthanum 

La 

138  2 

Tungsten         .... 

W 

184. 

Lead      .    .    . 

Pb 

2O6  QC. 

Uranium  .        .... 

u 

^  , 
2^9.6 

Li 

7.  02 

Vanadium  

V 

51-4 

Magnesium 

Me 

2A.  3 

Ytterbium  • 

Yb 

177. 

Manganese  

Mn 

C.C.. 

Yttrium   

Yt 

89.1 

Mercury  

Hg 

->D' 
2OO. 

Zinc  

Zn 

65.3 

Molybdenum 

Mo 

0.6 

Zirconium         .    . 

Zr 

9O.6 

TABLE  OF  THE  WEIGHTS  OF  GASES. 


TABLE  OF  THE  WEIGHTS  OF  GASES 

UNDER    STANDARD   CONDITIONS    (O°  C.    AND    760   MM.). 

Weight  of  i  Liter  in  Grams. 

Air  .  .       —  .  .  .  1.29327 

Oxygen  .                Oa  .  .  1-4295 

Nitrogen  .                N2  .  .  .  1.2572 

Hydrogen  .  .       HS  .  .  .           .0900 

To  find  the  weight  of  a  liter  of  any  other  gas,  multiply  its  specific  gravity 
compared  with  air  as  the  unit  by  the  weight  of  a  liter  of  air,  or  multiply  half 
the  molecular  weight  of  the  gas  by  the  weight  of  a  liter  of  hydrogen. 

Wt.  of  i  Liter  =  \  X  Mol.  Wt.  X  .0900. 


u6 


LABORATORY  WORK  IN   CHEMISTRY. 


TABLE  OF  TENSION  OF  AQUEOUS  VAPOR 

IN   MM.    OF  MERCURY. 


t°c. 

MM. 

t°c. 

MM. 

t°c. 

MM. 

—  10 

2.08 

II 

9-79 

40 

54.91 

-  9 

2.26 

12 

10.46 

45 

71.39 

-  8 

2.46 

13 

II.  16 

50 

91.98 

ij 

2.67 

14 

11.91 

55 

II7.48 

-  6 

2.89 

15 

12.70 

60 

148.79 

-  5 

3-13 

16 

13.54 

65 

186.94 

-  4 

3-39 

r7 

14.42 

70 

233.08 

-  3 

3-66 

18 

I5-36 

75 

288.50 

—    2 

3.96 

19 

16.35 

80 

354.62 

—    I 

4.27 

20 

17-39 

85 

433-00 

o 

4.60 

21 

18.50 

90 

525.39 

I 

4.94 

22 

19.66 

95 

633-69 

2 

5-30 

23 

20.89 

99 

733-21 

3 

b.69 

24 

22.18 

100 

760.00 

4 

6.10 

25 

23-55 

IOI 

787-59 

5 

6-53 

26 

24-99 

105 

906.41 

6 

7.00 

27 

26.51 

no 

1075.37 

7 

7-49 

28 

28.10 

8 

8.02 

29 

29.78 

9 

8.57 

30 

3L55 

10 

9.17 

35 

41.83 

INDEX. 


Acetylene 56 

Acids,  Bases,  and  Salts      ...  52 

Air,  Properties  of II 

Air,  Volumetric  Composition  of  .  23 

Aluminium 102 

Ammonia 39 

Ammonia,  Volumetric  Composition 

of 91 

Ammonium  Bromide,  Preparation 

of 65 

Ammonium    Compounds,    Reac- 
tions of 89 

Antimoniuretted  Hydrogen    .     .  80 

Antimony 79 

Apparatus,  Construction  of    .     .  I 

Aqueous  Vapor,  Tension  of  .     .  116 

Arsenates,  Reactions  of    ...  79 

Arsenic 77 

Arsine,  Arseniuretted  Hydrogen  78 
Atomic  Weights,  Table  of     .     .114 

Balance,  The  Chemical     ...  10 

Barium 94 

Bases,  Acids,  and  Salts     ...  52 

Bismuth 106 

Blowpipe,  The 61 

Boric  Acid 80 

Boron 80 

Bromides,  Reactions  of    ...  65 

Bromine 64 


Cadmium 97 

Calcium 91 

Calcium,  Reactions  of .     .     .     .  93 

Carbon 54 

Carbonates,  Reactions  of .     .     .  57 

Carbon  Dioxide 57 

Carbon  Dioxide,  Determination  of  93 

Carbon  Monoxide 59 

Chemical  and  Physical  Changes  .  3 

Chlorides,  Reactions  of     ...  52 

Chlorine 47 

Chromates 108 

Chromium 107 

Classification  of  Elements      .     .  83 

Cobalt 113 

Compounds,  Chemical,  defined    .  9 

Conservation  of  Matter     ...  20 

Copper 99 

Critical  Temperature    ....  49 

Definite  Proportions,  Law  of      .21 

Diffusion,  Law  of 30 

Dissociation,  Law  of    ....  92 

Elements  and  Compounds  defined     9 

Energy,  Conservation  of   .     .     .  20 

Equivalent  Weights  of  Metals    .  62 

Ethylene 56 

Fluorine    .                          ...  68 


117 


INDEX. 


Gas  Volumes,  Law  of  .     .     .     . 

PAGE 

46 

Gases,  Law  of  Diffusion  of    .     . 

30 

Gases,  Measurement  of    ... 

H 

Gases,  Weights  of  

"5 

Homogeneous  Matter  .... 

9 

Hydriodic  Acid  

67 

Hydrobromic  Acid  

64 

Hydrochloric  Acid  

49 

Hydrofluoric  Acid  

68 

Hydrofluosilicic  Acid  .... 

82 

Hydrogen      

26 

Hydrogen  Dioxide  

37 

Hydrogen  Sulphide      .... 

70 

Illuminating  Gas     

61 

Iodides,  Reactions  of  .... 

68 

Iodine  

66 

Iron     

in 

Kindling  Point,  The     .     .     .     . 

60 

Law  of  Conservation  of  Matter  . 

20 

Law  of  Definite  Proportions  .     . 

21 

Law  of  Diffusion  of  Gases     .     . 

30 

Law  of  Dissociation     .... 

92 

Law  of  Gas  Volumes    .... 

46 

Law  of  Multiple  Proportions 

37 

Law,  Periodic     

83 

Law  of  Reciprocal  Proportions  . 

62 

Law  of  Specific  Heats      .     . 

106 

Lead    

105 

Liquids,  Measurement  of  .     .     . 

10 

Liquefaction  of  Gases  .... 

49 

95 

no 

Manganese    .     ...     .     .     . 

109 

Marsh  Gas     

55 

Marsh's  Test  for  Arsenic  .     .     . 

78 

Matter,  Conservation  of    ... 

20 

Matter,  Homogeneous      .     .     . 

9 

Matter  and  Energy 7 

Measurement  of  Gases      ...  14 

Measurement  of  Liquids  .     .     .  10 

Mercury 97 

Mixtures,  Mechanical  ....  8 

Molecular  Weights 63 

Multiple  Proportions,  Law  of    .  37 

Nickel 113 

Nitrates,  Reaction  of   .     .     .     .43 

Nitric  Acid 40 

Nitric  Oxide 44 

Nitric  Oxide,  Volumetric  Compo- 
sition of 46 

Nitrogen 22 

Nitrous  Oxide 43 

Oxygen     .     . 16 

Ozone  . 36 

Periodic  Law 83 

Permanganates no 

Phosphates,  Reactions  of.  .  .  77 

Phosphine j6 

Phosphorus 75 

Physical  and  Chemical  Changes  .  3 

Potassium 83 

Potassium,  Reactions  of  .  .  .85 
Potassium  Bromide,  Preparation 

of 66 

Potassium  Chlorate,  Analysis  of .  21 
Potassium  Chlorate,  Preparation 

of 87 

Potassium  Iodide,  Preparation  of  86 

Rapidity  of  Chemical  Action      .  60 

Reciprocal  Proportions,  Law  of .  62 

Salts,  Acids,  and  Bases     ...  52 

Silicic  Acid 81 

Silicon      . 81 

Silicon  Hydride 8 1 

Silicon  Tetrafluoride    ....  82 


INDEX.  119 


PAGE 

Silver 100 

Silver  Nitrate,  Preparation  of     .  100 

Sodium 87 

Sodium  Chloride,  Preparation  of  88 

Specific  Gravity 10 

Specific  Heats,  Law  of     ...  106 

Spectroscope 89 

Stibine 80 

Strontium 94 

Sulphates,  Reactions  of    ...  75 

Sulphites,  Reactions  of     ...  72 

Sulphur 69 

Sulphur  Dioxide 71 

Sulphuric  Acid 73 


PAGE 

Table  of  Elements 1 14 

Table   of  Tension    of  Aqueous 

Vapor 116 

Table  of  Weights  of  Gases   .     .115 
Tin .104 

Volumes,  Law  of  Gas  ....     46 

Water,  Properties  of    ....     24 
Water,  Composition  by  Volume 

and  by  Weight      ....     32 

Zinc 96 

Zinc,  Equivalent  of      ....     62 


LD2l-lOOm-7,'40  (69368) 


